Textbook Chapter 12.9 - 12.16: Quantum Mechanics and Atomic Theory
How does electronic structure account for the properties of an element?
(from quantum number to chemical reactivity, from microscopic to macroscopic) 1) color of element in fireworks: transitions of electrons between energy levels (E = hν) 2) ionization energy (X -> X+ + e-) 3) electron affinities 4) atomic size 5) metals s nonmetals 6) magnetic properties (paramagnetic vs diamagnetic) 7) ionic charge 8) oxidation state 9) chemical reactivity
Why is the first ionization energy so much smaller than the second ionization energy?
After the first electron is removed, the resulting ion has a higher positive charge. This increased positive charge exerts a stronger attraction on the remaining electrons, making it harder to remove the second electron and thus raising the ionization energy. Additionally, the first electron is typically removed from a higher-energy orbital (such as a p orbital), while the second electron comes from a lower-energy, more tightly bound orbital (like an s orbital), which further explains the increase in ionization energy. A particularly large jump in ionization energy occurs when removing a core electron, as these electrons are much more tightly bound to the nucleus compared to valence electrons, leading to a significant increase in ionization energy.
Ionization example: aluminum (illustrative purposes)
Al (g) -> Al+ (g) + e- I1 = 580 kJ/mol Al+ (g) -> Al+2 (g) + e- I2 = 1815 kJ/mol Al+2 (g) -> Al+3 (g) + e- I3 = 2740 kJ/mol Al+3 (g) -> Al+4 (g) + e- I4 = 11,600 kJ/mol The first electron removed from the aluminum atom comes from the 3p orbital (Al has the electron configuration [Ne] 3s^2 3p^1) the second electron removed comes from the 3s orbital (since Al+ has the configuration [Ne] 3s^2)
Predict the trend in radius of the following ions: Be^+2, Mg^+2, Ca^+2, and Sr^+2
All these ions are formed by removing two electrons from an atom of a Group 2A element. In going from beryllium to strontium, we are going down the group, so the sizes increase. smallest Be^+2, Mg^+2 Ca^+2 largest Sr^+2
What is the periodic trend of reactivity and why does it occur?
As we go across a period from left to right, the reactivity of metals generally decreases, while the reactivity of nonmetals generally increases This occurs because, as we move across a period, the number of protons in the nucleus increases, which leads to a stronger nuclear charge. For metals, the stronger nuclear charge pulls the outer electrons closer to the nucleus, making it more difficult for them to be lost (since they have fewer energy levels and are more tightly bound). As a result, metals become less reactive because it becomes harder for them to lose electrons and form positive ions. For nonmetals, as the nuclear charge increases, the tendency to gain electrons becomes stronger because the outer electrons are more strongly attracted to the nucleus. This makes nonmetals more reactive as they more readily accept electrons to achieve a stable electron configuration
What is the periodic trend of electron affinity and why does it occur?
As we go across a period from left to right, the electron affinity generally increases it becomes more negative as you move from left to right This means that it gets more exothermic, and more energy is released when atoms gain electrons why? as the nuclear charge increases, the attraction between the nucleus and the added electron becomes stronger, leading to a more exothermic electron affinity More detailed explanations: As you move across a period, the number of protons (the nuclear charge) increases, which means the nucleus has a greater positive charge The added electron is attracted more strongly by the nucleus due to this increased nuclear charge, so it releases more energy when it gains it (makes the electron affinity more negative) ALSO Moving from left to right across a period, the atomic size decreases the outermost electrons are closer to the nucleus, increasing the attraction between the nucleus and the added electron the atom is more stable after gaining the electron, which results in a larger release of energy (more negative electron affinity).
What is the periodic trend of ionization energy and why does it occur?
As we go across a period from left to right, the first ionization energy generally increases This occurs because the number of protons in the nucleus increases, which leads to a stronger nuclear charge While core electrons shield outer electrons from this charge, electrons in the same principal energy level do not shield each other effectively As a result, the outer electrons experience a stronger attraction to the nucleus, making them harder to remove and thus increasing the ionization energy For example, note the sharp increase going across Period 2 from lithium through neon.
What is the periodic trend of ionic charge and why does it occur?
As we go across a period from left to right, the ionic charge of elements tends to become more negative (for nonmetals) or stay positive (for metals), with the charge becoming more pronounced for metals and less negative for nonmetals For metals (on the left side of the periodic table): The ionic charge tends to be positive and generally increases as you move from left to right in the metal part of the period For nonmetals (on the right side of the periodic table): The ionic charge tends to be negative and decreases (increasing in negativity) as you move across the period This occurs because the number of protons in the nucleus increases, which leads to a stronger nuclear charge. While core electrons shield outer electrons from this charge, electrons in the same principal energy level do not shield each other effectively. As a result, the outer electrons experience a stronger attraction to the nucleus, making it easier for metals to lose electrons (forming cations) and harder for nonmetals to lose electrons (forming anions). As you move from left to right across a period, nonmetals gain electrons to achieve a stable electron configuration, and their ionic charge becomes more negative. The more electronegative elements (further right) tend to form ions with greater negative charges, as they are more inclined to gain electrons to complete their valence shell.
What is the group trend of electron affinity and why does it occur?
As we go down a group from top to bottom, the electron affinity generally decreases becomes less negative (more positive) as you move down a group because the added electron is further from the nucleus, leading to weaker attraction and less energy release
What is the group trend of ionization energy and why does it occur?
As we go down a group from top to bottom, the first ionization energy generally decreases The main reason for the decrease in going down a group is that the electrons being removed are, on average, farther from the nucleus As n increases, the size of the orbital increases, and the electron is easier to remove, as they experience less nuclear pull (more shielded by the inner orbitals) For example, note the decrease in values for Group 1A and Group 8A.
What is the trend of reactivity down a group and why does it occur?
As we go down a group, the reactivity of metals generally increases, while the reactivity of nonmetals generally decreases. This occurs because the atomic size increases down a group, meaning that the outer electrons are farther from the nucleus and experience a weaker attraction to the positively charged nucleus. For metals, the larger atomic size means the outer electrons are more easily lost, leading to greater reactivity. As a result, alkali metals like lithium (Li) are less reactive than francium (Fr) at the bottom of the group, which loses its outer electron very easily. For nonmetals, the larger atomic size reduces the attraction of the nucleus for the outer electrons, making it harder for nonmetals to gain electrons. As a result, nonmetals like fluorine (F) are more reactive than iodine (I), as fluorine's smaller atomic radius allows it to attract electrons more strongly.
What was the Stern-Gerlach experiment?
Atom beam - splits into two beams from the original one beam Electrons those that sent up, ms of +½ spin up Electrons those that sent down, ms of -½ spin down Spin is quantized
How do we decide upon electron configuration?
Aufbau Principle Pauli exclusion principle Hund's Rule
What is the angular momentum quantum number?
represented by l related to SUBSHELL and SHAPE of orbital can have integral values from 0 to (n - 1) for each value of n relates to the angular momentum of an electron in a given orbital the dependence of the wave functions on l determines the SHAPE of the atomic orbitals the number of sub shells on the energy level is the different values of l (so it tells you what the maximum orbital kind it is) n = 2: Possible values of l are 0 and 1, indicating two subshells (2s and 2p) -- largest value of l is 1, so the largest orbital on that energy level is p n = 3: Possible values of l are 0, 1, and 2, indicating three subshells (3s, 3p, and 3d) -- largest value of l is 2, so the largest orbital on that energy level is d the TYPE/SHAPE of orbital is what this number depends on - specific number for each letter of orbital l = 0 corresponds to s l = 1 corresponds to p l = 2 corresponds to d l = 3 corresponds to f l = 4 corresponds to g and so on but they are PARTICULAR
What is the magnetic quantum number?
represented by ml related to ORIENTATION of orbital not as clear cut as the previous 2 as to if you see the number you know what is occurring ranges between numerical values of l and -l (including 0) relates to the orientation in space of the angular momentum associated with the orbital specific range for each letter of orbital - which the number then comes from s - ml range: 0 p - ml range: -1, 0, +1 d - ml range: -2, -1, 0, +1, +2 f - ml range: -3, -2, -1, 0, +1, +2, +3
What is the electron spin quantum number?
represented by ms the fourth quantum number property of ELECTRON, not ORBITAL accounts for emission spectra of atoms Electrons have two states, 'spin up' and 'spin down' In an orbital, each electron will be in a different spin state an orbital can hold 2 electrons because of spin up and spin down can only have one of two arbitrary values: +1/2 or -1/2 directly connected to Pauli Exclusion Principle
What is the principal quantum number?
represented by n can have integral values (1, 2, 3, 4...) related to SIZE and ENERGY of orbital The higher the n the higher the energy The higher the n, the electron spends more time further away from the nucleus As n increases, the orbital becomes larger and the electron spends more time farther from the nucleus An increase in n also means higher energy because the electron is less tightly bound to the nucleus, and the energy is less negative corresponds to ROW in periodic table
What is the order in which we fill electron shells?
row 1 1s row 2 2s, 2p row 3 3s, 3p row 4 4s, 3d, 4p row 5 5s, 4d, 5p row 6 6s, 5d^1, 4f, rest of 5d, 6p row 7 7s, 6d^1, 5f, rest of 6d, 7p
What are the four most common subshells?
s, p, d, f beyond f the letters become alphabetic: g, h, ... skipping j
What are the quantum numbers for the first four levels of orbitals in the hydrogen atom?
see chart
What does it mean to be diamagnetic?
slightly repelled by magnets and occur when all electrons are paired when exposed to an external magnetic field, these materials develop an induced magnetic moment that opposes the applied field, leading to a very weak repulsion from the magnetic field (if they even move) Materials like copper (Cu) and graphite are diamagnetic because all their electrons are paired. Ex: Silver Z = 47 (has unpaired electron) while iron Z = 26 has a lot of unpaired electrons - will be greatly attracted to the magnet while silver will not be as absorbed - fill seat
What are quantum numbers?
specify the properties of atomic orbitals and the properties of electrons in orbitals when boundary conditions are applied
What is different about the 2s and 3s orbital compared to the 1s orbital?
the 2s and 3s orbitals contain areas of high prob- ability separated by areas of zero probability These latter areas are called nodal surfaces, or simply nodes this is what we saw on the graph of probabilities when the ψ^2 crossed the x axis
What is electronic structure?
the arrangement and energy of electrons
What is electron affinity?
the energy change associated with the addition of an electron to a gaseous atom X (g) + e- -> X- (g) if the addition of the electron is exothermic, the corresponding value for electron affinity will carry a negative sign
What is hydration energy?
the energy released when water molecules surround and stabilize an ion
What is ionization energy?
the energy required to remove an electron from an atom or ion, where the atom or ion is assumed to be in ground state X (g) -> X+ (g) + e- given in terms of kJ/mole of atoms also common to see values given per atom --- in which case the term ionization potential is used (units: electron volts (eV) per atom, 1 eV = 1.602 x 10^-19 J)
What is the trend of ionic charge down a group and why does it occur?
the ionic charge of elements generally remains the same this is because all elements within a group have the same number of valence electrons, meaning they will typically form ions with the same charge when they gain or lose electrons to achieve a stable noble gas configuration
What is Koopman's Theorem?
the ionization energy is equal to the energy of the orbital from which it came it assumes that the electrons left behind in the resulting ion will not reorganize in response to the removal of an electron
Why are groups 1A, 2A, 3A, 4A, 5A, 6A, 7A, and 8A labeled the way they are?
the number indicates the total number of valence electrons for the atoms in these groups they are often called main-group, or representative elements ex: all the elements in group 5A (15) have the configuration ns^2 np^3, which is 5 valence electrons (the d electrons fill one period late and typically are not counted as valence electrons) 1A = group 1 2A = group 2 3A = group 13 4A = group 14 5A = group 15 6A = group 16 7A = group 17 8A = group 18 the ones skipped are the elements that fill d and f orbitals
What is charge density?
the ratio of the charge of an ion to its volume
What information is contained in the periodic table?
the similar properties of the atoms in a group rises from the identical valence electron configurations shared by group members- It is the number and type of valence electrons that primarily determine an atom's chemistry the electron configuration of any representative element - should memorize Cu and Cr tho bc of their strange configurations classifies elements into metals and nonmetals
How do you know how many orbitals are in a subshell?
you look at the value of the angular momentum quantum number (l) the number of orbitals in a subshell is calculated by the formula: 2l + 1 the number of possible orbitals is always one more than twice the value of l there is a specific number of orbitals per sub shell, which is why this works s = 1 p = 3 d = 5 f = 7 g = 9 and so on
What is a d orbital?
Butterfly shaped and can be oriented 5 ways labeled according to the axis of the Cartesian coordinate systems along which the lobes lie 10 electrons in a full set of d orbitals (xz, yz, xy, x^2-y^2, z^2) - since 5 orbitals, each orbital can hold 2e no 1d or 2d orbitals - begins in third row of periodic table (n = 3)
How can the dipping and increasing trend in ionization energies be explained?
By electron shielding and repulsions The increase in ionization energy from lithium to beryllium is expected because the 2s electrons do not shield each other effectively However, when moving from beryllium to boron, the ionization energy decreases because the 2s electrons shield the 2p electron effectively due to their closer proximity to the nucleus. From boron to nitrogen, ionization energy increases steadily because 2p electrons don't shield each other well The drop in ionization energy from nitrogen to oxygen is due to electron-electron repulsions in the doubly occupied 2p orbital of oxygen, making it easier to remove an electron Finally, from oxygen to neon, ionization energy increases as each electron is in a doubly occupied orbital, but the added electron repulsions in these orbitals cause differences in the ionization energies of oxygen, fluorine, and neon compared to boron, carbon, and nitrogen.
What are the properties of metals?
Chemical Property: Metals tend to lose electrons to form positive ions. Ionization Energy: Metals have low ionization energies. Location: Metals are primarily found on the left side of the periodic table. Reactivity: The most reactive metals are located in the lower-left corner of the table, where ionization energies are the smallest.
What are the properties of nonmetals?
Chemical Property: Nonmetals tend to gain electrons to form negative ions (anions) when reacting with metals. Ionization Energy: Nonmetals have large ionization energies and negative electron affinities. Location: Nonmetals are primarily found on the right side of the periodic table. Reactivity: The most reactive nonmetals are located in the upper-right corner, excluding the noble gases, which are very unreactive.
What transition metals are special?
Chromium and Copper Cr comes about as what should be [Ar] 3d^4 since it comes after V (3d^3) and before Mn (3d^5) HOWEVER it has a 4s^1 3d^5 configuration (so same number of electrons, just a strange configuration) - they do not know why Cu is the same - it comes about as what should be [Ar] 3d^9 since it comes after Ni (3d^8) and before Zn (3d^10) HOWEVER it has a 4s^1 3d^10 configuration (so same number of electrons, just a strange configuration) - they do not know why
How is the periodic table organized in terms of orbital filling?
Column 1 and 2 fill s orbitals Midsection fills 3d to 6d orbitals (and 4f and 5f) Column 13 to 18 fills 2p to 7p
What is a p orbital?
Dumbbell shaped and can be oriented 3 ways (x, y, z) has 2 lobes separated by a node at the nucleus per orbital -- has 3 orbitals so 6 lobes overall labeled according to the axis of the Cartesian coordinate systems along which the lobes lie ex: 2p orbital with lobes on x axis is the 2px orbital (2 is the n value, p is the l value corresponding sub shell, x is the orientation in space) 6 electrons in a full set of p orbitals (x, y, z) - since there is 3 orbitals (px, py, and pz) and each of them holds 2e each no 1p orbitals - begins in second row of periodic table (n = 2)
Where are the electrons found?
Each set of quantum numbers gives rise to a different wave equation that describes the waves that gives rise to the cloud visually of where these electrons are The probability equations - if we square the wave function (ψ), we get the probability of finding the electron in a given volume Eg: ψ^2 = 0.1 /pm^3 Normalized probability: (eg in volume of 2 pm^3) = (0.1/pm^3) x (2 pm^3) = 0.2 = 20% Electrons are not positioned exactly, as in Bohr atom, instead they are represented as clouds of probability In the Bohr atom, the distance was 0.529 A (where lowest electron was in distance from the center) - fixed orbit For schrodinger - the peak coincides with the bohr atom - radial probability (y) vs distance from nucleus (x)
How do we build atoms?
For given atom with given number of electrons, we have to build Fill from SMALLEST n levels first For hydrogen atom, only n determines energy levels For atoms with many electrons, n and other quantum numbers determine energy levels 1s level lowest Then 2s, 2p Then 3s, 3p - then 4s, level 4 is so spread out that 4s is in the 3 energy level, then 3d Then 4p, 4d, 4f Electrons fill all identical slots with one electron before putting two per spot
What do the first 20 elements look like in terms of electron affinities?
For the first 20 elements that form stable, isolated negative ions, all have negative (exothermic) electron affinities, meaning energy is released when an electron is added. The more negative the electron affinity, the greater the energy released.
Demo: put various salts of metals in dish + alcohol burning to give us flames - they are different colors
Lithium pink, sodium orange, potassium purple, strontium scarlet Different energy levels that are not evenly spaced - they have different delta E when you go up or down If we now heat the sample - sending electron up to higher level and when it falls back down, photons are emitted What is the energy of photons? Depends on distance between steps - the greater the space between the steps the greater the delta E Energy levels not evenly spaced filled orbitals to different degrees bc different metals - this is what results in different colors This is also the same metric by which fireworks work - different metals give rise to diff colors bc orbitals at different not evenly spaced energy levels Electrons get elevated and fall and emit photons
What are the properties of nonmetals?
Location: Metalloids are elements located along the division line between metals and nonmetals. Properties: Metalloids exhibit a mixture of both metallic and nonmetallic properties under certain conditions.
How do metals compare to nonmetals for ionization energy?
Staircase separates metals from nonmetals metal is easy to take electron away from - low ionization energy more difficult to remove electron from nonmetals than them - high ionization energy
What pattern develops on the periodic table for groups in terms of electrons?
The elements in the same group (vertical column of the periodic table) have the same valence electron configuration this is the reason why they have similar chemical properties (and were put in groups by Mendeleev) Elements with the same valence electron configuration often show similar chemical behavior (bonding) Only the principal quantum number of the occupied orbitals changes in going down a particular group
What relation does the number of nodes have to n?
The number of nodes increases as n increases For s orbitals, the number of nodes is given by n - 1 ex: 2s has 1 node
What about the noble gas electron configuration is special?
They have just the right number of electrons, as all the orbitals in all energy levels they occupy are full (all s, p, d, f in the level) Everyone wants to have a full energy level because it is the most stable way to be they are all at the very ends of their rows He - 1s^2 the only orbital is the 1s orbital, which is filled with 2e Ne - 1s^2 2s^2 2p^6 the 1s and 2s orbitals are filled as well as the 2p orbital - there are no other types of orbitals in these energy levels yet, so it is at max fullness pattern continues for every element in group 18 - noble gases
What does it mean to be paramagnetic?
attracted by a magnet and are created when unpaired electrons are present in the atom In the presence of an external magnetic field, these unpaired electrons align their spins with the field, causing the material to become magnetized in the direction of the field Example: Oxygen (O₂) is paramagnetic because it has two unpaired electrons in its molecular orbitals. if put in magnetic field and shot, it will separate into two populations
How can the exceptions be explained for electron affinity across a period?
electron configurations For example, nitrogen (with a half-filled p orbital, 1s²2s²2p³) resists adding an electron because doing so would result in a doubly occupied 2p orbital, increasing electron repulsion and making the ion unstable in contrast, carbon (with a 2p² configuration) can add an electron without such repulsion, forming a stable C⁻ ion.
How can the exceptions be explained for electron affinity across a group?
electron electron repulsions For example, in the halogens (Group 7A), while electron affinity decreases from chlorine to iodine, fluorine's electron affinity is anomalously lower than expected due to strong electron-electron repulsions in its small 2p orbitals. This makes the energy release for fluorine less than for chlorine, despite its high electronegativity.
What are core electrons?
electrons in lower energy levels that have already been filled ex: core electrons of nitrogen are the electrons in the 1s orbital because that energy level is full already and there are more electrons in more outer quantum levels
What is the Aufbau principle?
electrons occupy the orbitals of lowest energy first as protons are added one by one to the nucleus to build up the elements, electrons are similarly added to atomic orbitals electrons go into orbitals from low energy to high energy
What is the first ionization energy?
energy required to remove the highest-energy electron of an atom represented by I1
What is the second ionization energy?
energy required to remove the second highest-energy electron of an atom (or the most once the first is gone) represented by I2
What happens in row 6?
f orbitals begin lanthanides in row 6 - elements in which the 4f orbitals are being filled actinides in row 7 - elements in which the 5f orbitals are being filled
For principal quantum level n = 5, determine the number of subshells (differ- ent values of l) and give the designation of each.
for n = 5, the allowed values of l run from 0-4 l = 0 -> 5s l = 1 -> 5p l = 2 -> 5d l = 3 -> 5f l = 4 -> 5g
What are the special names for groups in the periodic table?
group 1: alkali metals (aka 1A) group 2: alkaline earth metals (aka 2A) groups 3-12: transition metals groups 13 (3A), 14 (4A), 15 (5A), and 16 (6A) have no other special name group 17: halogens (aka 7A) group 18: noble gases (aka 8A)
What does the function for the 2p orbital look like? (boundary surface representation)
has different signs in different regions of space ex: pz is positive where z is positive and is negative where z is negative the blue is the 90% probability of where the electrons are - Some probability of being way far out, but overall most are nearer to the center as you might expect from our discussion of the s orbitals, the 3p orbitals have a more complex probability distribution than that of the 2p orbitals -- but they can still be represented by the same boundary surface shapes The surfaces just grow larger as the value of n increases But there are nodes - spaces in volume where 0% probability Can take a piece of paper and slide in where the nucleus is between these two like oval jelly bean shaped areas where most are found
What happens after the transition metals?
implies we are already in period 4 at group 13, the p orbitals begin filling up (as the d orbitals are already filled) then the same trends continue until row 6 note: the (n + 1)s orbitals always fill before the nd orbitals ex: the 5s orbitals fill in rubidium and strontium before the 4d orbitals fill in the second row of transition metals (yttrium through cadmium)
What is the Pauli Exclusion Principle?
in a given atom, no two electrons can have the same set of four quantum numbers An orbital can hold only two electrons, and they must have opposite spins we know because of the Stern-Gerlach experiment IF TWO ELECTRONS SHARE AN ORBITAL / SET OF QUANTUM NUMBERS - they have to have either spin up and spin down (one of them has up and the other has down)
What are the ways an orbital can be illustrated?
in order of how they appear in the image: the square of the wave function probability distribution (most clear) - "slices" of the 3d electron density the surfaces that surround 90% of the total electron probability (the "sizes" of the orbitals)
What happens at Sc in the periodic table?
it begins a series of 10 elements called the transition metals (groups 3-12) their configurations are obtained by adding electrons to the five 3d orbitals (10 electrons needed to fill a 3d shell) so their configurations are written as [Ar] 4s^2 3d^(1 - 10)
What makes hydrogen different from the other members of group 1A?
it behaves as a nonmetal while the other members behave as a metal because of its very small size the electron in the small 1s orbital is bound very tightly to the nucleus
Consider atoms with the following electron configurations: 1s^2 2s^2 2p^6 1s^2 2s^2 2p^6 3s^1 1s^2 2s^2 2p^6 3s^2 Which atom has the largest first ionization energy, and which has the smallest second ionization energy? Explain your choices.
largest first ionization energy - 1s^2 2s^2 2p^6 since the 2p electrons do not shield each other very effectively, I1 will be large. The other configurations given include 3s electrons, which are effectively shielded by the core electrons and are farther from the nucleus than the 2p electrons in neon smallest second ionization energy - 1s^2 2s^2 2p^6 3s^2 For magnesium both I1 and I2 involve valence electrons For the atom with the configuration 1s^2 2s^2 2p^6 3s^1, the second electron lost is a core electron from a 2p orbital so it is holding it tightly - has to be #2
What is ml degeneracy?
means equal energy, asking how many spots of same energy level do we have so if we have 2p orbitals - 3 of them - no preference for which is filled first bc they are all at the same energy levels So like for a d orbital we have -2, -1, 0, 1, 2 so the value for ml is 5
What is ionic charge?
most common charge of an ion that the element would have (positive or negative or possibly just stable)
What does the function for the 1s orbital look like? (boundary surface representation)
positive everywhere in 3d space when evaluated at anywhere in space, it produces a positive number this is due to its spherical shape - no matter how you rotate it, the positivity/negativity stays the same
How many electrons are in an orbital?
2 electrons ex: there are 10 electrons in a d shell because there are 5 orbitals in a d shell and at 2 electrons per orbital that gets us 10 total
What is Hund's Rule?
2 parts 1) every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied 2) all electrons in singly occupied orbitals have the same spin (spins remain parallel) all 3 slots for an electron in the 2p orbital are equally likely to be filled For an atom with unfilled subshells, the lowest energy is achieved by electrons occupying separate orbitals, as allowed by the Pauli exclusion principle even though this is the case, there are 3 electrons in 3 separate 2p orbitals in a 1s^2 2s^2 2p^3 configuration -- it is written like they are all clumped though they are not
What does the function for the 3d orbital look like? (boundary surface representation)
4/5 ways - these 4 have 4 lobes centered in the plane indicated in the orbital label (like a clover!) dxz - centered in x plane dyz - centered in y plane dxy - centered in xy plane and lobes lie between the x and y axes d(x^2-y^2) - centered in xy plane and lobes lie along the x and y axes 1/5 - the fifth orbital dz^2 , has a unique shape with two lobes along the z axis and a "belt" centered in the xy plane the signs (phases) of the d orbital function are indicated inside the boundary surfaces
What is the electron configuration addition after lanthanum?
After lanthanum, which has the configuration [Xe] 6s^2 5d^1, a group of 14 elements called the lanthanide series, or the lanthanides, occurs This series of elements corresponds to the filling of the seven 4f orbitals (14 electrons) Note that sometimes one electron occupies a 5d instead of a 4f orbital This occurs because the energies of the 4f and 5d orbitals are very similar
Reactivity demo: 2M + 2H2O -> 2M+ + 2OHH- + H^2
Add water, metal charged, gives rise to OH- and H2 gas How to know it is reactive Bubbling hydrogen gas given off Have a pH indicator that turns a color as OH- released makes basic solution Lithium more active than magnesium Then sodium happened - metal melted, floating around - which was more reactive? In same column, lithium vs sodium - Na is more reactive - Mg is more reactive than Ca) More reactive down the column, further away from nucleus less shielding want to lose electrons
What is the electron configuration addition after actinium?
After actinium, which has the configuration [Rn] 7s^2 6d^1, a group of 14 elements called the actinide series, or the actinides, occurs This series corresponds to the filling of the seven 5f orbitals (14 electrons) Note that sometimes one or two electrons occupy the 6d orbitals instead of the 5f orbitals because these orbitals have very similar energies
What is an f orbital?
Firework shaped and can be oriented 7 ways 14 electrons in a full set of d orbitals - since 7 orbitals, each orbital can hold 2e no 1f or 2f or 3f orbitals - begins in fourth row of periodic table (n = 4) These orbitals are not involved in the bonding in any of the compounds we will consider in this text/exam Their shapes and labels are included here for com- pleteness. Because of their complexity, the phases of the f orbital functions are not represented in this diagram.
What is a valence electron?
an electron that is in the highest occupied principal quantum level of an atom ex: valence electrons of nitrogen are the electrons in the 2s and 2p orbitals because that energy level is not full yet most important electrons because they are involved in bonding
What elements on the periodic table are in which category?
Nonmetals: period 1- everything period 2- 4A-8A period 3- 5A-8A period 4- 6A-8A period 5- 7A-8A period 6- 8A period 7- 8A Metals: period 1- nothing period 2- 1A, 2A period 3- 1A-3A period 4- 1A-3A period 5- 1A-4A period 6- 1A-5A period 7- 1A-6A Metalloids: period 1- nothing period 2- 3A period 3- 4A period 4- 4A, 5A period 5- 5A, 6A period 6- 6A, 7A period 7- 7A
The first ionization energy for phosphorus is 1060 kJ/mol, and that for sulfur is 1005 kJ/mol. Why?
Phosphorus and sulfur are neighboring elements in Period 3 of the periodic table they have the following valence electron configurations: Phosphorus is 3s^2 3p^3, and sulfur is 3s^2 3p^4 Ordinarily, the first ionization energy increases as we go across a period, so we might expect sulfur to have a greater ionization energy than phosphorus However, in this case the fourth p electron in sulfur must be placed in an already occupied orbital The electron-electron repulsions that result cause this electron to be more easily removed than might be expected.
What is an s orbital?
SPHERE of electron density around the nucleus 2 electrons in a full set of 1 s orbital lower in energy than other orbitals because electron density is closest to the nucleus overall spherical shape becomes larger as n increases in each level - so there is 1s, 2s, 3s, 4s, 5s etc (begins in n = 1)
How does the electronic structure impact the ionization energy?
Successive ionizations Magnitude of first ionization energy Some ion in the gas phase - easiest to work with bc it is ions by themselves do not have to worry about how interacts with neighbors Pulling electron off - how much E required for that = ionization energy Sodium - E required to evict first electron is 495 kJ/mol - second electron almost 10x as much Magnesium - 735 - next one 1445, next one 7730 big jump The staircase showing us these jumps on periodic table Valence electrons at higher E levels farther from nucleus and core electrons close to nucleus - why is there this jump when pulling valence vs core electrons Sodium has 1 electron in 3s orbital - once you take it out, sodium resembles Neon and is happy (is essentially noble gas) - still demands energy bc trying to pull neg electron away from positive nucleus Taking the next electron is disrupting noble gas state Magnesium after taking 2 electrons away, it resembles noble gas configuration We can explain properties measured in lab Staircase separates metals from nonmetals - metal is easy to take electron away from - more difficult to remove electron from nonmetals than them Takes some amount of energy to pull off of hydrogen - Helium a lot more - plunges very far down to Li, then goes up a little more to Be (harder bc more nuclear charge holding it in), then gets a little easier (P orbital easier to take) - goes up with C then N, then goes down for oxygen (has to do with sharing seats - up until then, we have 3 electrons in seat - at oxygen we start sharing which is easier) If go down a column - it gets easier to take away the first electron nuclear charge is balanced by electrons - first electron you take off is furthest away - the further away from nucleus you get and shielded by inner shells, making it easier to pull it away
What is the periodic trend of atomic radius and why does it occur?
The atomic radius decreases as you move from left to right across a period This happens because the effective nuclear charge increases as more protons are added to the nucleus across a period, pulling the valence electrons closer to the nucleus and resulting in a smaller atomic size. increasing effective nu- clear charge (decreasing shielding) in going from left to right. This means that the valence electrons are drawn closer to the nucleus, decreasing the size of the atom.
What is the group trend of atomic radius and why does it occur?
The atomic radius increases as you move down a group it occurs because of the increases in the orbital sizes in successive principal quantum levels the additional energy levels and shells increase the distance between the nucleus and the outermost electrons, outweighing the effect of increased nuclear charge
How do fireworks work the way they do?
When fireworks explode, they heat up and emit light. The heat energy excites the electrons in the atoms or ions of the metal salts used in the fireworks. Electrons normally exist in specific energy levels (or orbitals) around the nucleus of an atom. When energy (in the form of heat) is supplied, such as during a firework explosion, it excites the electrons, causing them to move to a higher energy level or orbital. Once the electron has absorbed enough energy to jump to a higher orbital, it is not stable in that excited state. The electron will eventually "fall back" to its original lower-energy state. As it does, it releases energy in the form of light. The energy released when an electron drops back to a lower energy level corresponds to a specific wavelength of light. These wavelengths appear as bright lines in the emission spectrum, which are characteristic of the element or compound involved. Each element has a unique arrangement of energy levels, so the energy differences between these levels (and thus the light emitted when electrons transition between them) are unique as well. This results in a distinct bright-line spectrum for each element.
If map out the outermost orbital filled on periodic table, do you see patterns of electron cloud probabilities?
Yes columns fill orbitals (s, p, d, f) rows fill energy levels (n = 1, 2, 3...)
What does it mean to have degenerate orbitals?
a set of orbitals within an atom that have the same energy level essentially, they are energetically equal to each other so the 3 p orbitals within the n = 2 energy level are degenerate orbitals
What electron does ionization start with?
always the highest-energy electron (the one bound least tightly) that is removed first it is the most outer bound one (usually)
What are properties of the alkali metals?
among the most reactive of the metals INCREASE IN DENSITY going down group - typical occurs because atomic mass generally increases more rapidly than atomic size DECREASE IN MP AND BP - atypical (usually more complex relation) VERY REACTIVE low ionization energies and react readily with nonmetals to form ionic solids typically acts as the reducing agent in redox reactions Cs > Rb > K > Na > Li pattern of lower ionization energy being better reducer (true for solids but not liquids) for liquids, it is Li > K > Na due to the hydration energy of the ions in aqueous solution Li has much higher than K due to smaller size and higher charge density, meaning it is more easily hydrated and forms more readily in solution, despite its higher ionization energy while Li has the best reducing ability in aqueous solution, it does not react as violently with water as sodium or potassium - bc it has a high melting point, so it doesn't melt during the reaction, limiting its surface area in contact with water In contrast, Na and K have lower melting points and melt during the reaction, increasing their contact with water and making the reaction more vigorous
What are orbitals?
wave functions that satisfy the Schrödinger equation each is characterized by set of quantum numbers when boundary conditions applied
How can we make writing electron configurations easier?
we can abbreviate configurations using the noble gas configurations you put the noble gas symbol in brackets and that counts as all of the electron configuration that it has ex: sodium's electron configuration is 1s^2 2s^2 2p^6 3s^1 to avoid writing the inner level electrons (saving time), we can abbreviate this as [Ne] 3s^1 where the [Ne] represents the electron configuration of neon (1s^2 2s^2 2p^6)
What happens as we enter the third row of the periodic table?
we start with K since the 3p orbitals are fully occupied in argon, we might expect the next electron to go into a 3d orbital (recall that for n = 3 the orbitals are 3s, 3p, and 3d) BUT it is the 4s orbital instead since the 4 energy level shells are so spread out the chemistry of K is then very similar to that of Li and Na (indicating that the last electron is in the 4s rather than the 3d) then we have Ca, where the second electron comes to fill the 4s orbital then we have Sc
How do we write an electron configuration?
we write the energy level, then the shell letter, then the number of electrons in the orbitals within that shell this process is repeated for each type of energy level/shell/electrons combo ex: Boron has an electron configuration: 1s^2 2s^2 2p^1 if we look at the periodic table, we see that boron is in the 2nd row and the 13th column it retains all of the electrons of the elements before it (in order of fill order), so that is H, He, Li, and Be across rows, orbitals are filled - the 1s orbital is filled in row 1, and the 2s orbital is filled by Li and Be in row 2 Boron is the first element to come after these 4 elements, so it begins filling the next lowest energy orbital, which is 2p since each element across gains one more electron (to go with its one more proton), only one electron is added to the 2p orbital the 2p orbital can hold a total of 6 electrons, which will be filled further down the row but since boron only has 1 electron in the 2p orbital, that is where its configuration ends
