EK General Chemistry Lecture 1

Mass Number

# of protons & neutrons (with varying # of neutrons aka isotopes) Mass Number = Atomic Number - he atomic weight given on the periodic table can be assigned either of the commonly used units: atomic mass units (abbreviated as amu or the less common- ly used SI abbreviation, u) or grams/mole. Not every atom ofa given element will have amass equal to the mass number, because isotopes ofan atom have different numbers of neutrons. The atomic weight is equal to the weighted average of the naturally occurring isotopes of that element.

Atomic Number

# of protons (defines the element) * If the # of protons change, we'll have a different element. - An element may have any number of neutrons or electrons, but only one number of protons. - Protons and neutrons each have a mass of approximately 1 amu and the mass ofan atom is concentrated in the nucleus. This means that the mass number of an element is approximately equal to its atomic weight or molar mass.

Atomic Mass

# of protons + neutrons (e.g. Carbon¹²). Unit = amu

When is a atom electrically neutral?

# of protons = electrons

Chalcogens/Oxygen Group (Group 16)

*Oxygen & Sulfur are the important chalcogens.* 1. Oxygen (O) - O is the 2nd most electronegative element. - It is divalent & can form strong pi-bonds to make double bonds. - In nature, O exists as O₂ (dioxygen) and O₃ (ozone). - O typically reacts with metals to form metal oxides. - Alkali metals form peroxides (Na₂O₂) & super oxides (KO₂) with oxygen. 2. Sulfur - S₈: most common yellow solid form of pure sulfur - Metal sulfides (e.g. Na₂S) are the most common form of sulfur found in nature. - Sulfur can form up 6 bonds - Has the ability to pi-bond, forming strong double bonds.

Atomic Absorption Spectra

- "When an electron is excited to a higher energy level, it must absorb exactly the right amount of energy to make that transition." Excerpt From: Kaplan. "Kaplan MCAT General Chemistry Review: Created for MCAT 2015 (Kaplan Test Prep)." iBooks.

Dipole Moment

- A bond that has a dipole moment is POLAR. - A bond without a dipole moment is NONPOLAR. - dipole moment is represented by a vector pointing from the center of positive charge to the center of negative charge. The arrow is crossed at the center of positive charge, creating a plus sign. The dipole moment is measured in units of debye, D, and given by the equation: u=qd, where q is the magnitude of charge at either end of the dipole, and dis the distance between the centers of charge. * BE CAREFUL * - Symmetrical molecules will cancel out their dipole moments due to symmetry → leaving entire bond as a non polar bond (look at picture for CO₂).

Group 14

- All the Group 14 elements can form four covalent bonds with nonmetals. All beyond the second period can form two additional bonds with Lewis bases using d orbitals. Only carbon forms strong rr-bonds to make strong double and even triple bonds. This characteristic of carbon is often critical to the structure of biological molecules .

Induced dipole

- An induced dipole occurs when dipole moment is momentarily induced in an otherwise nonpolar molecule or bond by a polar molecule, ion, or electric field. The partial or full charge of the polar molecule or ion attracts or repels the elec- trons of the nonpolar molecule, separating the centers of positive and negative charge. Induced dipoles are common in nature and are generally weaker than permanent dipoles.

Instantaneous dipole moment

- An instantaneous dipole moment can arise spontaneously in an otherwise nonpolar molecule. Instantaneous dipoles occur because the electrons move about, and at any given moment they may not be distributed exactly between the two bonding atoms even when the atoms have equivalent electronegativity. Although instan- taneous dipoles are short-lived and weak, they can create an induced dipole in a neighboring molecule. - " The dipole moment of the polar bond or polar molecule is a vector quantity given by the equation: p = qd where p is the dipole moment, q is the magnitude of the charge, and d is the displacement vector separating the two partial charges. The dipole moment vector, represented by an arrow pointing from the positive to the negative charge, is measured in Debye units (coulomb-meters).

Noble Gases (Group 18)

- Are nonreactive, barely have any reactions → are very stable - Are all gases at room temperature

Radioactive Decay

- Atoms that spontaneously break apart. - All atoms (except Hydrogen) are subject to some type of spontaneous decay. - The mass & identity of an atom is located within its nucleus. - Atoms with a relatively high decay rate are RADIOACTIVE. * Radioactive Decay follows 1st order kinetics & is a type of EXPONENTIAL DECAY. *

Bond Energy

- Bond energy is the energy required to break a bond by separating its components into their isolated, gaseous atomic states. The greater the number of pairs of electrons shared between the atomic nuclei, the more energy is required to break the bonds holding the atoms together. Thus, triple bonds have the greatest bond energy, and single bonds have the lowest bond energy.

Halogens (Group 17)

- Consists of F, Cl, Br, I - Like to gain an electron to attain a noble gas configuration. - However, halogens other than "F" can take on oxidation states as high as +7 when bonding to other highly electronegative atoms such as oxygen. - When in compounds, "F" always has an oxidation state of -1 (means that F can make only 1 bond). - All of the halogens can combine with "H" to form gaseous hydrogen halides (which are then soluble in H2O, forming the hydrohalic acids). - Halogens react with metals to form ionic halides (e.g. NaCl). - Fluorine and chlorine are diatomic gases at room temperature; bromine, a diatomic liquid; and iodine, a diatomic solid. Halogens like to gain an electron to attain a noble gas configuration.

Group 15

- Group 15 elements can form 3 covalent bonds. In addition, all beyond the second period can form two additional covalent bonds by using their d orbitals. These elements can further bond with a Lewis base to form a sixth covalent bond. Nitrogen can form a fourth covalent bond by donating its lone pair of electrons to form a bond. Nitrogen forms strong rr-bonds to make double and triple bonds. Phosphorous can form only weak n-bonds to make double bonds. - The other Group 15 elements do not make n-bonds.

Paramagnetic VS. Diamagnetic Elements

- Hund's rule says that electrons will not fill any orbital in the same subshell until all orbitals in that subshell contain at least one electron, and that the unpaired electrons will have parallel spins. To understand the rule, consider that like charges repel each other. If we considered the energy of two particles with like charges, we would find that as the particles approach each other, the mutual repulsion creates an increase in potential energy. This is the case when electrons approach each other, so electrons avoid sharing an orbital when possible, spreading out amongst the orbitals ofa given subshell to minimize potential energy. - Hunds rule (low energy/high stability) creates 2 exceptions: 1. chromium (and other elements in its group) and copper (and other elements in its group). Chromium (Z = 24) should have the electron configuration [Ar] 4s23d4 according to the rules established earlier. However, moving one electron from the 4s subshell to the 3d subshell allows the 3d subshell to be half-filled: [Ar] 4s13d5 (remember that s subshells can hold two electrons and d subshells can hold ten). While moving the 4s electron up to the 3d-orbital is energetically unfavorable, the extra stability from making the 3d subshell half-filled outweighs that cost. 2. Similarly, copper (Z = 29) has the electron configuration [Ar] 4s13d10, rather than [Ar] 4s23d9; a full d subshell outweighs the cost of moving an electron out of the 4s subshell. Other elements in the same group have similar behavior, moving one electron from the highest s subshell to the highest d subshell. Similar shifts can be seen with f subshells, but they are never observed for the p subshell; the extra stability doesn't outweigh the cost. Paramagnetic elements: - Are attracted to a magnetic field - Unpaired e⁻'s in a shell (odd numbered e⁻) Diamagnetic elements: - Are not attracted to a magnetic field - All e⁻'s are paired (even numbered e⁻) BUT have to be careful just because all of the shells may not be filled. - The electron configurations of the transition metal ions are not the same as the nearest noble gas. For transition metals, ions are formed by losing electrons from the subshell with the highest principle quantum number first. Generally this is the s subshell. It is also important to know that there are a few exceptions to the electron configuration rules in the transition metals. Half-filled and filled subshells offer greater stability. Elements in Groups 6 and 11 are expected to have nearly half-filled or nearly filled d subshells. Instead they borrow one electron from the highest s subshell so they end up with a half-filled s subshell and a half-filled or filled d subshell.

COORDINATE COVALENT BONDS

- In a coordinate covalent bond, both of the shared electrons originated on the same atom. Generally, this means that a lone pair of one atom attacked another atom with an unhybridized p-orbital to form a bond, as shown in Figure 3.7. Once such a bond forms, however, it is indistinguishable from any other covalent bond. The distinction is only helpful for keeping track of the valence electrons and formal charges. Coordinate covalent bonds are typically found in Lewis acid-base reactions.

Naming inorgainc compounds

- Ionic compounds are named after their cation and anion. If the cation is transi- tion metal capable of having different charges, its name is followed by a Roman numeral in parentheses indicating the charge. Copper can take on a charge of 1+ or 2+, and is thus designated copper(!) ion or copper(II) ion. - Monatomic anions and simple polyatomic anions are given the suffix "-ide," such as hydride (H-) or hydroxide (OH-). Polyatomic anions with multiple oxygens end with the suffix "-ite" or "-ate," depending on the relative number ofoxygens. The more oxygenated species will use "-ate", such as nitrite ion (N02- ) versus nitrate ion (N03-). Ifthere are more possibilities, the prefixes "hypo-" and "per-" are used to indicate the fewest and most oxygens, respectively, such as hypochlorite (ClO-), chlorite (Cl02-), chlorate (Cl03- ), and perchlorate (Cl04- ) . If an oxyanion has a hydrogen, the word hydrogen is added, as in hydrogen carbonate ion (HC03-). The old name is for this ion is bicarbonate. - To name an ionic compound, put the cation name in front of the anion name, as in barium sulfate (BaS04) or sodium hydride (NaH). For binary molecular compounds (compounds with only two elements), the name begins with the name ofthe element that is farthest to the left and lowest in the periodic table. The name of the second element is given the suffix "-ide" and a number prefix is used for each element with more than one atom (e.g., dinitrogen tetroxide, N20 4). Acids are named for their anions. If the anion ends in "-ide," the acid name starts with "hydro-" and ends in "-ic," as in hydrosulfuric acid (H2S). With an oxyacid, the ending "-ic" is used for the species with more oxygens and "-ous" for the species with fewer oxygens, as in sulfuric acid (H2S04) and sulfurous acid (H2S03).

Metallic Bonds

- Metal + Metal bond

Hydrogen (H)

- Nonmetal → forms covalent bonds - It can also form ionic compounds with metal cations as the anion hydride. - Hydrogen plays a significant role in multiple chemical and physical processes, particularly acid-base chemistry and intermolecular forces. - it is an exception to the rule that elements in the same family have similar chemical properties.

The size of an atom has a significant effect on its chemistry.

- Small atoms hold charge in a concentrated way because they have fewer orbitals available to distribute and thereby stabilize charge. This concentra- tion of charge makes the smallest element in each group bonds more readily and with greater bond strength, especially when in ionic form. Fluorine is a good example. The fluoride ion (F) is too small to manage its full negative charge. Therefore, it is generally insoluble and bonds immediately when in solution. For this reason, in toothpaste, fluoride (a poison in high concentra- tions) bonds immediately with the enamel of teeth before it can be ingested. - Small atoms do not have d-orbitals available for bond formation, and there- fore cannot form more than 4 bonds. Large atoms have d-orbitals, allowing for more than 4 bonds. Oxygen typically forms two bonds, while the Larger sulfur can form up to six. Smaller atoms have the advantage when it comes to n-bonding. The p-orbitals on large atoms do not overlap significantly, so they cannot easily form n-bonds. Carbon, nitrogen, and oxygen are small enough to form strong n-bonds while their Larger third row family members form only weak n-bonds, if they form n-bonds at all.

Energy level of electrons

- The Aufbau principle (sometimes called the "building up principle") states that with each new proton added to create a new element, the new electron that is added to maintain neutrality will occupy the lowest energy level available. All other things being equal, the lower the energy state of a system, the more stable the system. Thus, electrons look for an orbital in the lowest energy level whenever they are added to an atom. The orbital with the lowest energy will be located in the sub- shell with the lowest energy. - This rule states that the lower the sum of the values of the first and second quantum numbers (n + l), the lower the energy of the subshell. This is a helpful rule to remember for Test Day. If two subshells possess the same (n + l) value, the subshell with the lower n value has a lower energy and will fill with electrons first. - For the representative elements, the shell level of the most recently added electrons is given by the period in the periodic table. The most recently added electron ofF is in shell 2, and the most recently added electron for Sr is in shell 5. For transition metals the shell of the most recently added electron lags one behind the period. The most recently added electron of Ag is in shell 4, though it is in the fifth period. For the lanthanides and actinides, the shell ofthe most recently added electron lags two behind the period. The most recently added electron for Ce is in shell 4, though the lanthanide series corresponds to the sixth period. - Valence electrons, the electrons which contribute most to an element's chemical properties, are located in the outermost shell of an atom. In most cases, only electrons from the s and p subshells are considered valence electrons.

Application of Bohr model

- The Bohr model of the hydrogen atom (and other one-electron systems, such as He+ and Li2+) is useful for explaining the atomic emission and absorption spectra of atoms. - As electrons go from a lower energy level to a higher energy level, they get AHED: Absorb light Higher potential Excited Distant (from the nucleus)

Atoms

- The radius of a nucleus is on the order of 10-4 angstroms (A). One angstrom is 10-10 m. The nucleus consists ofprotons and neutrons, collectively called nucleons. - protons and neutrons are approximately equal in size and mass.

When does a zero-rrder reaction occur?

- This occurs in enzyme-catalyzed reactions when the concentration of the substrate far OUTWEIGHS the concentration of the enzyme. - All of the enzyme catalytic sites get saturated with substrate & the addition of more substrate has no effect on reaction rate.

Transition Metals

- When the transition metals form ions, they lose electrons first from their highest s-subshell and then from their d-subshell,

Zeff Trend

- Zeff generally increases going left to right across the periodic table. While more protons are added across a period, increasing Z, the new electrons added are in Electron Salty roughly the same energy level, and therefore do not experience significantly more shielding than the previous electron. - Zeff also increases going from top to bottom down the periodic table. Though the energy level of the outermost electrons increases down a group, the attractive pull of the growing positively charged nucleus outweighs the additional shielding effects of higher electron shells. - Zeff drops going from neon to sodium. This happens because the new electron is added to an entirely new shell, the 3s subshell. This causes a strong increase in shielding and reduction in Zeff, but the outermost electron in sodium still experiences a higher Zeff than the outermost electron of the element immediately above it on the periodic table, lithium. This is because the effects of the more strongly charged nucleus outweigh the shielding effect that an additional electron shell can provide. A similar drop also occurred between He and Li, though it was not quite as large because there were fewer protons and electrons involved.

Coulomb's law, F = kq1q2/r^2

- coulomb's law describes the electrostatic force between an electron and the nucleus. Electrostatic force is the force between charged objects, which is attractive between opposite charges and repulsive between like charges. - the distance between the electron and the nucleus is r. For q1 we might plug in the positive charge ofthe nucleus, Z, and for q2, the charge on an electron, e. - this would work well for hydrogen, where the lone electron feels 100% of the positive charge on the nucleus. However, in helium the first electron shields some of the nuclear charge from the second electron, so that it doesn't feel the entire nuclear charge Z. - Shielding occurs due to the repulsive forces between electrons. The amount of charge felt by the most recently added electron is called the effective nuclear charge (Zeff)· - In complete shielding, each electron added to an atom would be completely shielded from the attractive force of all the protons except for the last proton added, and the Zeffwould be 1 eV (electron volts) for each electron. Without shielding, each electron added would feel the full attractive force ofall the protons in the nucleus, and the Zeff would simply be equal to Z for each electron. - The Zeff• and not Z, should be plugged in for q1 in Coulomb's law to find the force on the outermost electron. The force on an electron is a function of both q1 (Zeff) and r (the distance from the nucleus). Notice that Zeff can be used to explain the atomic radius trend discussed earlier. Since the effective nuclear charge increases from left to right on the periodic table, each additional electron is pulled more strongly toward the nucleus. The result is that atoms tend to get smaller when adding electrons across the periodic table. When moving down a group, each drop represents the addition of a new electron shell, and thus atoms tend to increase in size moving down a group. - Zeff can also be used to understand trends in ion size. As discussed previously, cations are smaller than the neutral element while anions are larger. When an atom loses an electron, Zeff increases because there are now more protons relative to electrons. Increased Zeff means electrons are pulled closer to the nucleus. When an atom gains an electron, Zeff decreases because there are now more electrons relative to protons. Shielding increases due to the increased repulsive forces between electrons, and electrons are pushed further away from the nucleus.

Metalloids

- have some metallic and some non-metallic characteristics.

Octet rule (exceptions)

- octet rule, which states that an atom tends to bond with other atoms so that it has eight electrons in its outermost shell, thereby forming a stable electron configuration similar to that of the noble gases. - A simple way to remember all the exceptions is as follows: 1. Incomplete octet: These elements are stable with fewer than eight electrons in their valence shell and include hydrogen (stable with 2 electrons), helium (2), lithium (2), beryllium (4), and boron (6). 2. Expanded octet: Any element in period 3 and greater can hold more than eight electrons, including phosphorus (10), sulfur (12), chlorine (14), and many others. 3. Odd numbers of electrons: Any molecule with an odd number of valence electrons cannot distribute those electrons to give eight to each atom; for example, nitric oxide (NO) has eleven valence electrons.

Isoelectronic ions

- soelectronic ions are ions with the same number of electrons, but different elemental identities. 0 2- , F , neutral Ne, Na+, and Mg2+ all have the same number of electrons. However, since they have different numbers of protons, the electrons feel different effective nuclear charges. As a result, isoelectronic ions are not the same size. The largest of these is 0 2- . Its electrons feel the weakest Zeff because there are two more electrons than protons. The smallest of these is Mg2+. Its electrons feel the strongest Zeff because there are two more protons than electrons.

Quantum Numbers

- the first quantum number is the principle quantum number, n. It designates the shell level of the electron, with low numbers closest to the nucleus. Remember, a larger integer value for the principal quantum number indicates a larger radius and higher energy. This is similar to gravitational potential energy. - The second quantum number is the azimuthal quantum number, l. It designates the electron's subshell, each of which has a distinct shape. l= 0 is the s subshell, l= 1 is the p subshell, l= 2 is the d subshell, and l= 3 is the f subshell. The azimuthal quantum number is very important because it has important implications for chemical bonding and bond angles. A simpler way to remember this relationship is that the n-value also tells you the number of possible subshells. Therefore, there's only one subshell (l = 0) in the first principal energy level; there are two subshells (l = 0 and 1) within the second principal energy level; there are three subshells (l = 0, 1, and 2) within the third principal energy level, and so on. **For any principal quantum number n, there will be n possible values for l, ranging from 0 to (n - 1). The energies of the subshells increase with increasing l value; however, the energies of subshells from different principal energy levels "may overlap. For example, the 4s subshell will have a lower energy than the 3d subshell. - The third quantum number is the magnetic quantum number, mL, which designates a precise orbital within a given subshell. Each orbital can hold two electrons. Each subshell has orbitals with magnetic quantum numbers ranging from -L to +L. Within the p-subshell (L = 1), there are three magnetic quantum numbers: -1, 0, and 1. An atom will spread out its electrons amongst the orbitals, such that all orbitals in a subshell have one electron before any has two. The number of total orbitals within a shell is equal to n^2. Solving for the number of orbitals for each shell gives 1, 4, 9, 16... Since there are two electrons in each orbital, the number of elements in the periods of the periodic table is 2, 8, 18, and 32. For any value of l, there will be 2l + 1 possible values for ml. For any n, this produces n2 orbitals. For any value of n, there will be a maximum of 2n2 electrons (two per orbital). - The fourth quantum number is the electron spin quantum number, ms. The electron spin quantum number has possible values of -1/2 or +1/2. This final quantum number is used to distinguish between the two electrons that may occupy the same orbital, since those electrons will have the same first three quantum numbers. - Let's summarize: The flrst quantum number is the shell. It corresponds roughly to the energy (eve( of the electrons within that shell. The second quantum number is the subshell. It gives the shape. Rec- ognize that s orbitals are spherical and porbitals are dumbbell-shaped. The third quantum number'is the specific orbital within a subshell. The fourth quantum number distin- guishes between two electrons in the same orbital; one is spin +1/2 and the other is spin -1/2.

Periodic Trends

- these predictions can be summarized as four periodic trends: atomic radius, ionization energy, electronegativity, and electron affinity. As a general rule, the atomic radius increases going "down" and "to the left" on the periodic table, while the other properties increase "up" and "to the right." - Atomic radius is the only periodic trend that increases moving down a group and across a period from right to left. The largest atomic radius in the Periodic Table belongs to cesium (Cs, 260 pm), and the smallest belongs to helium (He, 25 pm). - Know how ionic radii differ from atomic radii - trends.

Electron Configuration Background

1. As the electrons get further away from the nucleus, the energy level of electrons increases. - Electrons in higher shells are at a higher energy level (electrostatic energy increases). - Be certain that the total number of electrons in your electron configuration equals the total number of electrons in the atom. Notice that an electron can momentarily absorb energy and jump to a higher energy level, creating an atom in an excited state. -

Shell filling pattter

1. Atoms lose electrons from the highest energy shell first. This means that in transition metals, electrons are lost from the s-subshell first, and then from the d-subshell. 2 . Ions seek symmetry. Representative elements form noble gas electron configurations when they make ions. For example, group 1 atoms form 1+ cations and group 17 atoms form 1- anions. Transition metals try to "even- out" their d-orbitals so that each orbital has the same number of electrons. Whenever possible, an ion will have a half-filled or completely filled orbital. Important reference points in the periodic table include the periods 1 and 2, which have half-filled and completely filled sorbitals, periods 7 (VIIB) and 12 (liB), which have half-filled and completely filled d orbitals, and periods 15 and 18, which have half-filled and completely filled porbitals.

Element

1. Each atom in the universe can be identified as belonging to a particular element. 2. Elements are the building blocks of compounds & cannot be broken down into simpler substances by chemical means.

Nonpolar Covalent Bond vs. Polar Covalent Bond

1. Nonpolar Covalent Bond - When e⁻'s are shared equally by 2 atoms w/ equal electronegativity. 2. Polar Covalent Bond - When e⁻'s are NOT shared equally b/c of a difference in electronegativity.

Protons, Neutrons, & Electrons

1. Protons have a positive (+) charge, while neutrons are electrically neutral. 2a. Protons & neutrons are approx. equal in size and mass (approx. 1 amu). 2b. Protons & neutrons are held together to form the nucleus by the strong nuclear force. 3. Electrons have a negative (-) charge & surround the nucleus & are much SMALLER than neutrons & protons. Surrounding the nucleus at a distance of about 1 to 3 A are electrons. The mass ofan electron is more than 1800 times smaller than the mass ofa nucleon. - The mass of an electron is approximately that of a proton. Because subatomic particles' masses are so small, the electrostatic force of attraction between the unlike charges of the proton and electron is far greater than the gravitational force of attraction based on their respective masses. **LOCATION** - Protons & neutrons are found in the nucleus. - Electrons are found in shells or orbitals that surround the nucleus of an atom. - Since the nucleons are tiny compared to the distance between the nucleus and the outermost electrons, the atom itself is composed mostly of empty space. - A charge of 1 e is equal to 1.6 x 10-19 coulombs, the SI unit for charge. An atom is electrically neutral when it contains the same number ofprotons and electrons. - A mole is a number of "things" (atoms, ions, molecules) equal to Avogadro's number, NA = 6.02 × 1023. For example, the atomic weight of carbon is which means that the average carbon atom has a mass of 12.0 amu (carbon-12 is far more abundant than carbon-13 or carbon-14), and 6.02 × 1023 carbon atoms have a combined mass of 12.0 grams.

Hydrogen Bond

1. The STRONGEST type of dipole-dipole interaction - (strongest type of INTERmolecular force). H-bond is responsible for high BP of H2O. 2. BE CAREFUL - Although hydrogen bonds are the strongest INTERmolecular forces, it is still much weaker than any covalent bond (a INTRAmolecular force). ***Hydrogens loves their FON (phone) → Fluorine, Oxygen, Nitrogen. *** (F-H, O-H, N-H)

Binding Energy

1. The energy that is required to break the nucleus into separate protons & neutrons. 2. The stability of the nucleus can be measured by its binding energy.

Vapor Pressure (Gas)

1. VP is the measure of how many molecules escape into the gas phase. - If in the liquid phase, the molecules have greater intermolecular forces, then its going to be harder for molecules to escape into the gas phase. ** SO, HIGHER INTERMOLECULAR FORCES LEADS TO A LOWER VAPOR PRESSURE.**

Van der Waals Forces or london dispersion force

1. WEAK dipole-dipole forces (NON-polar bonds) - All molecules exhibit Van der Waal forces, even when they are capable of stronger intermolecular interactions. 2. Surface Area - The more surface area you have, the more london dispersion forces you'll have. TEMPORARY/TRANSIENT WEAK DIPOLE

What happens when a neutral atom loses an electron? What happens when a neutral atom gains an electron?

1. When a neutral atom loses an electron, it becomes a cation (+) & its size gets smaller. 2. When a neutral atom gains an electron, it becomes an anion (-) & its gets bigger.

Gamma Decay

A gamma ray is a high frequency photon. It has no mass or charge, and does not change the identity of the atom from which it is given off. Gamma decay/gamma ray emission occurs when an electron and positron collide.

Catalyst

A substance that INCREASES the RATE of a reaction without being consumed or permanently altered. - Increase the rate of both the forward & reverse reactions. - Enhance product selectivity & reduce energy consumption.

Alkali Metals (Group 1)

Alkali Metals characteristics: - Low densities & low melting points - Easily form 1+ cations, such as Na+. - Highly reactive (reacting with most nonmetals to form ionic compounds). - React with "H" to form hydrides (e.g. NaH). - Highly reactive w/ H2O →react exothermically (explosively) w/ H2O to produce metal hydroxide & hydrogen gas (H₂). *In nature, alkali metals exist only in compounds.* - The alkali metals have only one loosely bound electron in their outermost shells. Their Zeff values are very low, giving them the largest atomic radii of all the elements in their respective periods. This low Zeff value also explains the other trends: low ionization energies, low electron affinities, and low electronegativities. Alkali metals easily lose one electron to form univalent cations, and they react readily with nonmetals—especially the halogens—as in NaCl.

Alkaline Earth Metals (Group 2)

Alkaline Earth Metal characteristics: - Harder, more dense, & melt at higher temperatures than alkali metals. - Forms ²⁺ cations (e.g. Mg²⁺) - Form basic solutions with H2O - Less reactive than alkali metals because their highest energy electron completes the s orbital. - Heavier alkaline earth metals are more reactive than lighter ones. *Alkaline earth metals exist only in compounds in nature.*

Electron Configuration Exceptions

All subshells have to be "filled" or at least "half-filled." These exceptions are for: Cr, Mo (in Group 6) ; Cu, Ag, Au (in Group 11). Look in notebook to see thorough explanation.

Halides

An atom that is bonded to a halogen (F, Cl, BR, I)

Intermediates

Are species that are products of 1-step & reactants of a later step in a multistep reaction. - Because they get used up before the end of the reaction, they are not shown in the overall chemical equation. - Intermediates are often present only in low concentrations.

Atomic Emission Spectra

At room temperature, the majority of atoms in a sample are in the ground state. However, electrons can be excited to higher energy levels by heat or other energy forms to yield excited states. Because the lifetime of an excited state is brief, the electrons will return rapidly to the ground state, resulting in the emission of discrete amounts of energy in the form of photons. - The electromagnetic energy of these photons can be determined using the following equation: E = hc/lambda ; c is the speed of light and lambda is the wavelength. - These energy transitions do not form a continuum, but rather are quantized to certain values. Thus, the spectrum is composed of light at specified frequencies. It is sometimes called a line spectrum, where each line on the emission spectrum corresponds to a specific electron transition. Because each element can have its electrons excited to a different set of distinct energy levels, each possesses a unique atomic emission spectrum, which can be used as a fingerprint for the element. - The Bohr model of the hydrogen atom explained the atomic emission spectrum of hydrogen, which is the simplest emission spectrum among all the elements. The group of hydrogen emission lines corresponding to transitions from energy levels n ≥ 2 to n = 1 is known as the Lyman series. The group corresponding to transitions from energy levels n ≥ 3 to n = 2 is known as the Balmer series, and includes four wavelengths in the visible region. The Lyman series includes larger energy transitions than the Balmer series; it therefore has shorter photon wavelengths in the UV region of the electromagnetic spectrum. The Paschen series corresponds to transitions from n ≥ 4 to n = 3. - The energy associated with a change in the principal quantum number from a higher initial value ni to a lower final value nf is equal to the energy of the photon predicted by Planck's quantum theory. Combining Bohr's and Planck's calculations, we can derive: E = hf/lambda = -Rh[1/ni^2 - 1/nf^2]. - Note that unlike other equations, this is initial minus final; the negative sign in the equation accounts for absorption and emission. Thus, a positive E corresponds to emission, and a negative E corresponds to absorption. - This is expressed best by the Heisenberg uncertainty principle: It is impossible to simultaneously determine, with perfect accuracy, the momentum and the position of an electron.

Why do ions matter?

Because charge= reactivity and that's where the is. Also, biological systems rely on ions to regulate cellular activities, as in the formation of an action potential in a neuron.

Mnemonic for Catalysis for Rate Laws (Includes Intermediates)

Catalysts is CRAP → CRP In a chemical equation, catalyst appears first as a REACTANT then a PRODUCT. Intermediates - In a chemical equation, intermediates appear first as a PRODUCT then a REACTANT (reverse of a catalyst).

Cations vs. Anions

Cations are positive ions (more protons, less electrons). Anions are negative ions (more electrons, less protons). Example: A salt is a compound composed of a positive & a negative ion together. ** Changing the number of neutrons creates an isotope; changing the number of electrons creates an ion; changing the number of protons changes the atom into a different element.

Covalent Bonds

Covalent Bond characteristics: - Nonmetal + Nonmetal bond (e.g. HCl) - Nonmetals would rather gain e⁻'s than lose so they SHARE e⁻'s - Covalent bonds are INTRAmolecular force (bonds w/in molecules). - Low MP & BP (e.g. boiling H2O) - Forms molecules which are made up of INTERmolecular bonds

Reaction Rate

Describes how quickly the concentration of the reactants or products are changing over the course of the reaction.

What happens to concentration of the decaying substance?

During radioactive decay, the concentration of the decaying substance, decreases EXPONENTIALLY.

Valence Electrons

Electrons in the outermost shell. - Elements in the same group on the periodic table have similar chemical properties b/c they have the same # of valence electrons. S holds 2 electrons P holds 6 electrons D holds 10 electrons F holds 14 electrons Shapes of s and p orbitals. - Elements in the sarne group on the periodic table have similar chemical properties because they have the same number of valence electrons, or electrons in the outermost shell. They tend to make the same number ofbonds and exist as similarly charged ions. - The electrons closer to the nucleus are at lower energy levels, while those that are further out (in higher shells) have higher energy. The electrons that are farthest from the nucleus have the strongest interactions with the surrounding environment and the weakest interactions with the nucleus. These electrons are called valence electrons; they are much more likely to become involved in bonds with other atoms because they experience the least electrostatic pull from their own nucleus. - The sharing of these valence electrons in covalent bonds allows elements to fill their highest energy level to increase stability.

Electron Affinity

Energy change associated with gaining an additional electron. More precisely, it is the energy released (exothermic) when an electron is added to an isolated atom. - Just like electronegativity, electron affinity tends to increase on the periodic table from left to right and from bottom to top. The sign of electron affinity values can be different for different atoms because some atoms release energy when accepting an electron (and thus become more stable), while others require energy input to force the addition of an electron (since the additional electron decreases stability). 1. Electron affinity is more exothermic (releasing heat) to the right and up on the periodic table. 2. Electron affinity values for the noble gases (group 18) are endothermic (absorbs heat), because noble gases are stable, so significant amounts of energy are required to force them to take on electrons and become less stable.

Emission Line Spectrum

Excited e⁻ drops from HIGH to LOW energy → energy is RELEASED. EMITS a photon.

Absorption Line Spectrum

E⁻ goes from LOW to HIGH energy state → energy is absorbed from photon/light. - Planck theorized that electromagnetic energy is quantized, that it comes only in discrete units related to the wave frequency. If energy is transferred from one point to another via an electromagnetic wave, and we wish to increase the amount of energy transferred, the energy can only change in discrete increments given by: dE =hf where h is Planck's constant (6.6 X 10-34 J s) and f is frequency. - instein showed that if light is considered as a particle phenomenon, where each photon is one particle, the energy of a single photon is given by the same equation: dE photon = hf. - Neils Bohr applied the quantized energy theory to create the Bohr model, where electrons rotate around the nucleus on a path characterized by a certain energy level (shell). His model explained the line spectra for hydrogen but failed for atoms with more than one electron. Louis de Broglie expanded the model to demonstrate the wave characteristics of electrons and other moving masses with this equation: lambda =h/mv where h is Planck's constant, m is mass, and v is velocity. - Atoms have both particle and wave properties; therefore atomic structure cannot be understood by considering electrons only as moving particles.

When are INTERmolecular forces broken?

First off, intermolecular forces are forces between separate molecules. Intermolecular forces are broken when a compound moves from the liquid to gas phase (a physical change).

What happens to the identity of an element after it undergoes radioactive decay?

First... - Recall that the elemental identity of an atom is defined by the # of PROTONS. - If the # of protons CHANGE (which happens in some types of radioactive decay) → change in the IDENTITY of an element to a NEW element. Secondly... - ADDITION of a proton → a move to the RIGHT of the periodic table to determine the NEW element. - LOSS of a proton → move to the LEFT.

Relationship between Concentration & Time for 0, 1st, 2nd, 3rd Order Reactions. (Picture → Reaction Order Graphs)

For the MCAT, memorize the graphs for 0 & 1st order rate laws. - Note that the relationships shown are linear, with slopes equal to the rate constant for the given rate law. * A typical question might draw your attention to the descending green line & ask you if the value of the rate constant decreases over time. - ANSWER: IT DOES NOT! → the rate constant is CONSTANT! (hence the name) .

HIGHER Intermolecular Forces Summary

HIGHER INTERMOLECULAR FORCES = - HIGHER MP - HIGHER BP - HIGHER VISCOSITY - HIGHER SURFACE TENSION - LOWER VAPOR PRESSURE *****

3 Types of Radioactive Decay

I. alpha decay (α) II. beta decay (β) III. gamma decay/ production of gamma rays (γ)

Intermolecular forces and its effect on MP & BP

INTERmolecuar forces affect melting point (MP) & boiling point (BP), viscosity, surface tension & vapor pressure. Melting Point Example: 1a. When you melt a solid, the solid is held together in a rigid crystal structure and some of the intermolecular forces that are holding it together from molecular compounds. 1b. When you break some of those structures, to get the solid into a liquid. The stronger the intermolecular forces, the more energy it'll take to break those, and the higher the MP. Boiling Point Example: 2a. In a liquid, the molecules are in contact but they're moving. 2b. BUT, in a gas, the molecules are separated by huge amounts of space. So, you have to break all of the remaining intermolecular forces to boil a liquid. **THE STRONGER THE INTERMOLECULAR FORCES, THE HARDER ITS GOING TO BE TO BOIL → THER HIGHER THE BP.**

Stoichiometry

Involves figuring out the quantities of products and reactants in a chemical equation.

Ionic Bonds

Ionic Bond characteristics: - Metal + Nonmetal bond (e.g. NaCl) e⁻'s are transferred from one atom to another. - These bonds are a special type of electrostatic interaction. - INTRAmolecular force - High MP & BP (e..g trying to melt salt requires high temp). - Brittle (forms crystals as in NaCl) ***IONIC BONDS DO NOT FORM MOLECULES, INSTED THEY FORM CRYSTALS SO THEY DO NOT HAVE ANY INTERMOLECULAR FORCES***.

Isoelectronic Ions

Ions with the same number of electrons, but different elemental identities. Example: O²⁻, F⁻, Na⁺, Mg²⁺ - All have the same number of electrons. However, since they have different numbers of protons, the electrons feel different effective nuclear charges. - ∴ isoelectronic ions are not the same size.

Limiting Reactant/Reagent

Is the REACTANT that would be completely used up if the reaction were to run to completion.

Theoretical Yield

Is the amount of PRODUCT we expect to be created when a reaction runs to completion. *The amount of product created by a real experiment is the ACTUAL YIELD.*

Nuclear Decay

It includes the degradation of particles in the nucleus of an atom. * RECALL THAT... * - Atomic nuclei are held together by strong nuclear force. - Without this strong nuclear force, protons would repel ("fight off") one another. - Neutrons SPACE out the protons to prevent this, in order to stabilize the nucleus.

What does a "reaction runs to completion" mean?

It means that a reaction generates product until the supply of at least one reactant is fully finished.

What happens when a molecule is bulky (sterically hindered)?

It will be more difficult for other molecules to react with it → slow step in a reaction.

Kinetic vs. Thermodynamic

Many reactions have several possible products, each favored by different reaction conditions. THERMODYNAMIC - The thermodynamic product is MORE stable, but is produced more SLOWLY because the required energy input is HIGHER. - HIGH temperatures drive the reaction towards the thermodynamic product. KINETIC - The kinetic product is LESS stable but can be produced more FASTER because the required energy input is LOWER. - LOW temperatures favor the kinetic product.

Metals & Electrons

Metal atoms have a loose hold on their outer electrons. ∴ electrons move easily from one metal to the next, transferring energy or charge in the form of heat or electricity.

Metals

Metals are large atoms that tend to lose electrons to form positive ions (cations) & positive oxidation states. - Metals can be described as atoms in a sea of electrons, which emphasizes their loose hold on their electrons and the fluid-like nature of their valence electrons. The easy movement of electrons within metals is what gives them their metallic character. Metallic characteristics: - Luster - Malleability (easily hammered into thin strips) - Ductility (easily stretched) - Conductive (Heat & Electricity) - electrons move easily from one metal atom to the next, transferring energy or charge in the form of heat or electricity. - HIGH BP & MP (exception is Mercury, which has a lower MP than room temp).

Molecular Formula vs. Empirical Formula

Molecular Formula: - Gives you the number of atoms in every single molecule (e.g. Benzene → C₆H₆) Empirical Formula: - Is the most reduced whole number ratio formula (e.g. Benzene → CH - a 1:1 ratio according to the molecular formula) * When finding the empirical formula from the percent mass composition, ALWAYS pretend that the sample weighs exactly 100g. - The empirical formula and the atomic weight of each element can be used to calculate the percent composition by mass for a compound. To find percent composition, multiply an element's atomic weight by the number of atoms it contributes to the empirical formula. Divide the result by the net weight ofall the atoms in the empirical formula, which yields the mass fraction of that element in the compound. Multiply by 100% to get the percent composition by mass.

7 Diatomic Molecules to Remember

Never Have Fear Of Ice Cold Beer Never - Nitrogen (N₂) Have - Hydrogen (H₂) Fear - Fluorine (F₂) Of - Oxygen (O₂) Ice - Iodine (I₂) Cold - Chlorine (C₂) Beer- Bromine (B₂) Example: The statement "Nitrogen is nonreactive" refers to N₂ (diatomic) instead of N.

Nonmetals

Nonmetal characteristics: - Form COVALENT BONDS with one another & becomes MOLECULES. - Have LOWER MP than metals - Tend to form anions (-) → commonly react with metal cations (+) to form ionic compounds. - They have high ionization energies, electron affinities, and electronegativities, as well as small atomic radii and large ionic radii. They are usually poor conductors of heat and electricity. All of these characteristics are manifestations of the inability of nonmetals to easily give up electrons. Nonmetals are less unified in their chemical and physical properties than the metals.

What does this chemical reaction tell us? N₂(g) + 3H₂(g) = 2NH₃(g)

N₂(g) + 3H₂(g) = 2NH₃(g) Details: - The equation does not tell us anything about what we have. - Instead, it tells us the ratio in which things react. - It tells us that - for every 1 mole of N₂ that we do have, it requires 3 H₂'s to react with which will form 2 moles of NH₃. - 1:3:2 ratio

If the difference in electronegativity is significant, the bond is said to have ?

Partial Ionic Character

Periodic Table

Period: each horizontal row (across). Groups/Families: the vertical (down) columns. *Memorize the following groups*: Alkali metals (1), Alkaline earth metals (2), Oxygen group (16), Halogens (17), and Noble gases (18).

Semi-log Plot

Plotting the logarithm of amount of atoms as a function of time would produce a straight line semi-log plot.

Reaction Orders

Reaction orders indicate how changes in the reactant concentrations influence the reaction rate.

Low Energy State

Remember: 2 atoms will ONLY form a bond if they can lower their overall energy level b/c nature seeks the lowest possible energy state. ∴, energy is released when bonds are formed → lowering their overall energy.

Representative Elements

Representative elements form ions with the electron configurations of noble gases (Group 18) in order to gain stability. - the representative elements make ions by forming the closest noble gas electron configuration; this is why metals tend to form cations and nonmetals tend to form anions.

Bohr Model

Represents the atom as a nucleus surrounded by electrons in separate electron shells. * This model explained the line spectra for hydrogen but failed for atoms with more than 1 electron. - Bohr used Planck's quantum theory to correct certain assumptions that classical physics made about the pathways of electrons (planck relation: E=hf, h is proportionality constant or planck's constant 6.626x10-34 Jxs and f is frequency of the radiation). Classical mechanics postulates that an object revolving in a circle, such as an electron, may assume an infinite number of values for its radius and velocity. The angular momentum (L = mvr) and kinetic energy (k=1/2 mv^2) of the object could therefore take on any value. However, by incorporating Planck's quantum theory into his model, Bohr placed restrictions on the possible values of the angular momentum. Bohr predicted that the possible values for the angular momentum of an electron orbiting a hydrogen nucleus could be given by: L = nh/2pi. - Bohr then related the permitted angular momentum values to the energy of the electron to obtain: E= - R(h)/n^2, where RH is the experimentally determined Rydberg unit of energy, equal to Therefore, like angular momentum, the energy of the electron changes in discrete amounts with respect to the quantum number. hTerefore, the electron in any of its quantized states in the atom will have an attractive force toward the proton; this is represented by the negative sign. - Ultimately, the only thing the energy equation is saying is that the energy of an electron increases—becomes less negative—the farther out from the nucleus that it is located (increasing n). Energy (E) is directly proportional to the principal quantum number (n) (due to - sign) - ground state (all systems like this unless subjected to extreme temp or irradiation) vs excited state

Isotopes

Same # of protons but different # of neutrons. Isotopes have similar chemical properties (e.g. ¹²C, ¹³C, ¹⁴C → all have the same number of protons but different neutrons).

Quantum Mechanics

States that elementary particles can only lose or gain energy in discrete small units (e.g. walking up stairs instead of walking up a ramp). Chad's explanation: There's only energies you're allowed to have - only certain quantities of energy you're allowed to have.

Pauli Exclusion Principle

States that no 2 electrons in the same atom can have the same 4 quantum numbers. - The first quantum number is the principle quantum number. - There are three magnetic quantum numbers: -1, 0, and 1.

Heisenberg Uncertainty Principle

States that there you can never simultaneously know the exact position and the exact momentum (speed) of an object. - this uncertainty arises from the dual nature (wave-particle) ofmatter, and is on the order ofPlanck's constant (6.6 x 10-34 J s): dxdp >/= h/2 Example: The more we know about the position of a particle, the less we know about the momentum.

2nd Requirement for a Given Collision to Initiate a Reaction

The atoms of both molecules must ALIGN in a specific way for the collision to result in a reaction. - When the molecules do not align correctly → NO REACTION OCCURS. - Even if the particles have sufficient KE to overcome the Ea. ** BOTH criteria must be met for a collision to initiate a reaction. **

Beta Decay (β-decay)

The breakdown of a neutron into a proton & electron, & the expulsion of the newly created electron. - Since a neutron is destroyed, but a proton is created, the mass number stays the same, but the atomic number increases by one.

Bond Length

The distance between the nuclei of 2 atoms in a bond when they are at their lowest possible energy state. - For a given pair of atoms, a triple bond is shorter than a double bond, which is shorter than a single bond. *If the atoms are pushed too close or pulled too far apart than their bond length, they will be at a higher energy level.*

Bond Dissociation Energy (Bond Energy)

The energy needed to separate a bond completely.

2nd Ionization Energy

The energy required to remove a 2nd electron from the SAME atom to form a +2 cation. **The 2nd ionization energy is ALWAYS GREATER than the first because once 1 electron is removed, the effective nuclear charge increases for the remaining electrons.**

1st Ionization Energy

The energy required to remove an electron from a neutral atom in its gaseous state to form a +1 cation.

Ionization Energy

The energy required to remove an electron from an atom. - Ionization energy generally increases along the periodic table from left to right and from bottom to top. - the ionization energy trend can be remembered by considering its relation to zeff and r. Moving across a period to the right, increasing zeff values pull electrons more strongly toward the nucleus. Therefore, more energy is required to rip them off Moving down a group, z .lf increases, but the distance of the electron from the nucleus increases as well. Coulomb's law demonstrates that the electrostatic force, F, decreases with the square ofthe distance from the nucleus, r. Due to the exponent, the increased distance plays a more important role than the increased Zelf and the attractive electrostatic force decreases down a group. Less force means less energy is required to remove the electron, so ionization energy decreases moving down a group. Example: When an electron is more strongly attracted to the nucleus, more energy is required to remove it (harder to break it up). - ** remember - Remember: 1) pulling the outermost electron closer decreases atomic radius; 2) holding the outermost electron more tightly increases ionization energy; 3) atoms with greater z.ff will pull more strongly on electrons in covalent bonds, increasing electronegativity across a period; and 4) atoms with stronger zeff will more readily accept another electron, so electron affinity increases across a period. - Removing an electron from an atom always requires an input of heat, which makes it an endothermic process.

Electrostatic Force

The force between charged objects - Attractive between OPPOSITE charges (-) & (+) - Repulsive between LIKE charges (-) & (-)

Half-Life

The length of time necessary for one half of a given amount of a substance to decay.

Alpha Decay (α-decay)

The loss of an α particle. - An α particle is a helium nucleus, meaning it contains 2 protons & 2 neutrons.

Rate Law

The rate law can be used to determine how changes in initial concentration affect the reaction rate.

Rate-Determining Step

The rate of the SLOWEST elementary step determines the rate of the OVERALL reaction. EXAMPLE: - If the 1st step is the slow step, the rate law is derived from this step.

1st Requirement for a Given Collision to Initiate a Reaction

The relative KE's of the colliding compounds must be GREATER THAN or EQUAL to a threshold energy called the activation energy (Ea). - In a given sample of a compound, each particle is moving at a different speed - only the particles with sufficient speed will have the KE needed to overcome the Ea.

Kinetics

The study of reaction rates and mechanisms. - Deals with the rate of a reaction as it moves TOWARDS EQUILIBRIUM.

Electronegativity

The tendency of an atom to attract electrons shared in a covalent bond. - When 2 atoms have different electronegativities, they share electrons unequally, which causes polarity (e.g. H2O). - Like ionization energy, electronegativity tends to increase across a period from left to right and up a group. The most commonly used measurement of electronegativity is the Pauling scale, which ranges from a value of 0.79 for cesium to a value of 4.0 for fluorine. It is important to remember that fluorine is the most electronegative element. **Fluorine is the MOST electronegative element.** - The electronegativity of hydrogen falls between that of boron and that of carbon. When bonded with hydrogen, carbon and elements to the right of carbon will carry a partialnegative charge while hydrogen will carrya partial positive charge. Think of CH4 • Boron and the elements to the left of boron will carry a partial positive charge when bonded to hydrogen, while the hydrogen will carry a partial negative charge. Thinkofthe hydrides. (H) in NaH or LiAIH3 . - Atoms with large differences in electronegativity (1.6 or larger on the Pauling scale as a rule of thumb) will form ionic bonds. Metals and non-metals usually exhibit large electronegativity differences and form ionic bonds with each other. Atoms with moderate differences in electronegativities (0.5 - 1.5 on the Pauling scale) will generally form polar covalent bonds. Atoms with very minor electronegativity differences (0.4 or smaller on the Pauling scale) will form nonpolar covalent bonds.

What does Thermodynamics deal with?

Thermodynamics deals with the balance of reactants & products AFTER they have achieved equilibrium.

Arrhenius Equation

This equations tells us that as Ea INCREASES, the rate constant DECREASES. The rate constant is affected by PRESSURE, CATALYSTS & TEMPERATURE. - Higher pressure INCREASES rate constant - Catalysts lower the Ea, thus INCREASING the rate constant. - As the temperature increases, the number of collisions increases (thus increasing the molecules capability to overcome the Ea. → THIS MEANS: the rate constant increases with increases temperature. - Higher temperatures increase the rate of both the forward and reverse reactions → making it EASIER to reach equilibrium FASTER.

Photoelectric Effect

Was used by Einstein to demonstrate the particle aspect of light. a. To show that light was made of particles called photon. Einstein's Theory: Light is made up of particles. b. The electron must be ejected one-to-one photon-electron collisions instead of combined energies of 2 or more photons.

What happens when activation energy is lowered?

When Ea is LOWERED, MORE collisions have enough KE to result in a reaction. - This leads to MORE reactions & an increase in the overall reaction rate. SUMMARY: The reaction rate depends exponentially on the Ea.

Chemical Reactions

When a compound undergoes a reaction and changes its bonding or structure to form a NEW compound. Examples: - Combustion - Metathesis - Redox Chemical reactions are represented by chemical equations, with the reactants & products: e.g. CH₄ + 2O₂ ⇌ CO₂ + 2H₂O

Physical Reactions

When a product undergoes a reaction & maintains its molecular structure (its IDENTITY remains the same). Examples: - Melting - Evaporation - Dissolution - Rotation of Polarized Light

Valence e⁻ Exception

When elements are written in their "ground-state configurations", we do not count filled "d-subshells" as valence e⁻'s. Once the d-subshell is full (w/ 10 e⁻'s) → it is part of a previously filled shell (look in notebook for examples. - Note that all systems tend toward minimal energy; thus on the MCAT, atoms of any element will generally exist in the ground state unless subjected to extremely high temperatures or irradiation.

Electron Config. "equations"

n =# of shell n² = # of orbitals 2n² = # of max. e⁻'s

Ion

number of electrons ≠ number of protons → the atom will carry a charge.