CHEM TEST 3

The formation of crystalline NaCl lattice from sodium cations and chloride anions is highly exothermic and more than compensates for the endothermicity of the electron transfer process. In other words, the formation of ionic compounds is not exothermic because sodium "wants" to lose electrons, rather it is exothermic because of the large amount of heat released when sodium and chloride ions coalesce to form a crystal lattice.

Why is the formation of solid sodium chloride from a solid sodium and gaseous chlorine exothermic even though it takes more energy to form the Na+ ion than the amount of energy released upon formation of Cl-?

•penetration: when an electron penetrates the electron cloud of the 1s orbital and experiences the charge of the nucleus more fully because it is less shielded •as outer electron undergoes penetration into the region occupied by inner electrons, experiences greater nuclear charge and, according to Coulomb's law, a lower energy

What is penetration? How does the penetration of an orbital into the region occupied by core elements affect the energy of an electron in that orbital?

•shielding: one electron is blocked from full effects of the nuclear charge so that the electron experiences only part of the nuclear charge •Inner (core) electrons shield outer electrons from full nuclear charge

What is shielding? In an atom, which electrons tend to do the most shielding (core or valence)?

-metals conduct electricity bc e- are free to move -flow of e- in response to e- produces electric current -good conductors bc very mobile e- -no specific bonds in a metal, so can be deformed easily and let ions slide and form new shape

how does the electron sea model explain the conductivity of metals? the malleability and ductility of metals?

-covalent bonds are highly directional-attraction is due to sharing of e- -each bond links one specific pair of atoms -ionic: nondirectional and hold an array of ions-fundamental units of covalently bonded compounds are individual molecules -interactions are more weak-when it melts/boils, molecules remain intact -only weak bonds need to be overcome so the melting/boiling points are lower

how does the lewis model for covalent bonding account for the relatively low melting and boiling points of molecular compounds (compared to ionic compounds)

combinations of atoms that satisfy the octet rule on each atom are stable, and those that don't aren't stable

how does the lewis model for covalent bonding account for why certain combinations of atoms are stable while others are not

single bond: one pair of electrons shared between 2 atoms- double bond: 2 e- pairs are shared between the same 2 atoms. shorter and stronger than single bonds triple bond: 3 e- pairs shared between 2 atoms. even shorter and stronger

in what ways are double and triple covalent bonds different from single covalent bonds

-ability of an atom to attract e- to itself in a chemical bond -results in a polar bond -increases across a period -decreases down a column -most electronegative is F

what is electronegativity? what are the periodic trends in electronegativity

-fictitious charge assigned to each atom in a lewis structure to distinguish competing structures -charge it would have if all bonding e- were shared equally between bonded atoms -number of valence e- - (lone pairs + 1/2 bonding e-) -distinguish between competing skeletal structures or competing resonance structures

what is formal charge? how is formal charge calculated? how is it helpful?

-write correct skeletal structure for molecule -add up number of valence e- of each atom - this is number of e- total in the structure -distribute e- among atoms, giving octets for as many atoms as possible -if an atom lacks an octet, make double or triple bonds to give them

what is the basic procedure for writing a covalent lewis structure?

when metal atoms bond together to form a solid, each metal atom donates one or more e- to an electron sea

what is the electron sea model for bonding in metals?

1.6x10^-19 x 100 pm /10^12 pm /3.34 x 10^-30 C x m = 4.8 D

what is the magnitude of the dipole moment formed by separating a proton and an electron by 100 pm? 200 pm?

-not all atoms have 8 e- surrounding them -odd octets: odd number of e- ex. NO -incomplete octets: fewer than 8 e- ex. BF3 -expanded octet: more than 8 e- ex. AsF5

why does the octet rule have exceptions? give 3 major categories of exceptions and an example of each

-bonding pair: pair of electrons shared between two atoms -lone pair: pair of atoms associated with only one atom and not involved in bonding

within a covalent lewis structure, what is the difference between lone pair and bonding pair electrons?

The modern periodic tables is credited primarily to the russian chemist Dmitri Mendeleev. Mendeleev's table is based on the periodic law, which states that when elements are arranged in order of increasing mass, their properties recur periodically. Mendeleev arranged the elements in a table in which mass increased from left to right and elements with similar properties fell in the same columns.

Who is credited with arranging the periodic table? How are the elements arranged in the modern periodic table?

•sub levels for multi electron atoms split because penetration of the outer electrons into the region of the core electrons •sub levels are not split because they are empty in the ground state

Why are the sub levels within a principle level split into different energies for multi electron atoms but not for the hydrogen atom?

Chemical bonds form because they LOWER THE POTENTIAL ENERGY between the charged particles that compose the atom. Bonds involve the attraction and repulsion of charged particles.

Why do chemical bonds form? What basic forces are involved in bonding?

The rows of the periodic table grow progressively longer because you are adding sub levels as the n level increases.

Why do the rows in the periodic table get progressively longer as you move down the table? For example, the first row contains two elements, the second and third rows each contain 8 elements, and the fourth and fifth rows each contain 18 elements. Explain.

The degree of mixing between two orbitals decreases with increasing energy difference between them. Mixing of the 2s and 2px orbitals is greater in B₂, C₂, and N₂ than in O₂, F₂, and Ne₂ because in B, C, and N, the energy levels of the atomic orbitals are more closely spaced in the O, F, and Ne. This mixing produces a change in energy ordering for the pi₂p and the sigma₂p molecular orbitals.

Why does the energy ordering of the molecular orbitals of the period 2 diatomic molecules change in going from N2 to O2?

•Pauli exclusion principle: no two electrons in an atom can have the same four quantum numbers .•two electrons in same orbital have three identical quantum numbers (n, l, ml), they must have different spin •Pauli exclusion principle implies each orbital has a max of two electrons with opposing spins

Why is electron spin important when writing electron configurations? Explain in terms of the Pauli exclusion principle.

According to VESPR theory, the repulsion between electron groups on interior atoms of a molecule determines the geometry of the molecule.

According to VSEPR theory, what determines the geometry of a molecule?

In Lewis theory, a chemical bond is the sharing or transferring of electrons to attain stable electron configurations for the bonding for the bonding atoms. If electrons are transferred, the bond if an ionic bond. If the electrons are shared, the bond is covalent bond.

According to the Lewis model, what is a chemical bond?

Main group: •Atomic radius decreases across period (left to right). -Adding electrons to same valence shell -Effective nuclear charge increases -Valence shell held closer •Atomic radius increases down group. -Valence shell farther from nucleus -Effective nuclear charge fairly close Atomic radius is an average radius of an atom based on measuring large numbers of elements and compounds. A more general term, in the atomic radius, refers to a set of average bonding radii determined from measurements on a large number of elements and compounds. The atomic radius represents the radius of an atom when it is bonded to another atom and is always smaller than the van der Waals radius. a. As you move right across a period in a periodic table, atomic radius decreases b. As you move down a column in the periodic table, atomic radius increases

Define atomic radius. For main-group elements, give the observed trends in atomic radius as you... a. Move across a period in the periodic table. b. Move down a column in the periodic table.

When writing the ele config of a transition metal cation, remove the ele's in the highest n-value orbitals first, even if this does not correspond to the reverse order of filling. Normally, even though the d orbital ele's add after the s orbital ele's, the s orbital ele's are lost first. This is because: 1)the ns and (n-1)d orbitals are extremely close in energy and, depending on the exact config, can vary in relative energy ordering 2)As the (n-1)d orbitals begin filling in the first transition series, the increasing nuclear charge stabilizes the (n-1)d orbitals relative to the ns orbitals. This happens BC the (n-1)d orbitals are not outermost orbitals and therefore are not effectively shielded from the increasing nuclear charge by the ns orbitals.

Describe how to write an electron configuration for a transition metal cation. Is the order of electron removal upon ionization simply the reverse of electron addition upon filling? Why or why not?

Bonds are formed when atoms attain a stable electron configuration. Because the stable configuration usually has eight electrons in the outermost shell, this is known as the octet rule.

Describe the octet rule in the Lewis model.

a. Cations are much smaller than their corresponding parent. This is BC the outermost ele's are shielded from the nuclear charge in the atom and contribute greatly to the size of the atom. When these ele's are removed to form the cation, the same nuclear charge is now acting only on the core ele's. b. Anions are much larger than their corresponding atoms. This is because the extra ele's are added to the outermost ele's but no additional protons are added to increase the nuclear charge. The extra ele's increase the repulsions among the outermost ele's, resulting in an anion that is larger than the atom.

Describe the relationship between: a. the radius of a cation and that of the atom from which forms b. the radius of an anion and that of the atom from which it forms

The IE2 of Mg involves removing the 2nd outermost ele's leading to an ion with a noble gas config for the core ele's. The IE3 requires removing a core ele from an ion with a noble gas config. This requires a tremendous amount of energy, making IE, very high. For Al, IE3 involves removing the 3rd outermost ele for Al, leaving the ion with a noble gas config of the core ele's, IE4 then requires removing a core ele from an ion with a noble gas config. This requires a tremendous amount of energy and make IE4 very high.

Examination of the first few successive ionization energies for a given element usually reveals a large jump between two ionization energies. For example, the successive ionization energies of magnesium show a large jump between IE2 and IE3. The successive ionization energies of aluminum show a large jump between IE3 and IE4. Explain why these jumps occur and how you might predict them.

Locate the element on the periodic table. Identity it's group number to determine the amount of valence electrons and the row number to determine the highest principal quantum number.

Explain how to write the electron configuration for an element based on its position in the periodic table?

A paramagnetic species has unpaired electrons in one or more molecular orbitals. A paramagnetic species is attracted to a magnetic field. The magnetic property is a direct result of the unpaired electron(s). The spin and angular momentum of the electrons generate tiny magnetic fields. A diamagnetic species has all of its electrons paired. The magnetic fields caused by the electron spin and orbital angular momentum tend to cancel each other. A diamagnetic species is not attracted to a magnetic field and is, in fact, slightly repelled.

Explain the difference between a paramagnetic species and a diamagnetic one.

The electron geometry is the geometrical arrangement of the electron groups around the central atom.The molecular geometry is the geometrical arrangement of the atoms around the central atom.The electron geometry and the molecular geometry are the same when every electron group bonds two atoms together. The presence of unbonded lone pair electrons gives a different molecular geometry and electron geometry.

Explain the difference between electron geometry and molecular geometry. Under what circumstances are they not the same?

In valence bond theory, hybrid orbitals are weighted linear sums of the valence atomic orbitals of a particular atom, and the hybrid orbitals remain localized on that atom.In molecular orbital theory, the molecular orbitals are weighted linear sums of the valence atomic orbitals of all the atoms in a molecule, and many of the molecular orbitals are delocalized over the entire molecule.

Explain the difference between hybrid atomic orbitals in valence bond theory and LCAO molecular orbitals in molecular orbital theory.

The lettered group number of a main-group element is equal to the number of valence electrons for that element.

Explain the relationship between a main-group element's lettered group number (the number of electron's column) and its valence electrons.

The chemical properties of elements are largely determined by the number of valence electrons they contain. Their properties are periodic because the number of valence electrons is periodic. Because elements within a column in the periodic table have the same number of valence electrons, they also have similar chemical properties.

Explain the relationship between properties of an element and the number of valence electrons that it contains.

For main group elements the row number is the highest principal quantum number in its electron configuration. For transition elements the principal quantum number is the row number minus 1. In inner transition elements it is the row number minus 2.

Explain the relationship between the element's row number in the periodic table and the highest principal quantum number in the element's electron configuration. How does this relationship differ from main-group elements, transition elements and inner transition elements?

The simple answer is that the s block can contain 2 electrons, the p block 6 electrons, the d block 10 electrons and the f block 14 electrons. Because of this, there are 14 different outer electron configurations and hence it spans 14 groups. Only 2 orbitals with 2 in each is needed to fill the elements in the s block while the elements in the p block need 6 each.

Explain why the s block in the periodic table has only two columns while the p block has six.

-by adding up the number of valence electrons in each atom -same just add or subtract according to charge

how do you determine the number of electrons that go into the lewis structure of a molecule? a polyatomic ion?

•Atoms in the same group increase in size down the column. •Atomic radii of transition metals are roughly the same size across the d block. -Much less difference than across main-group elements -Valence shell ns^2, not the (n −1)d electrons -Effective nuclear charge on the ns^2 electrons approximately the same a. The radii of transition elements stay roughly constant across each row instead of decreasing in size as in the main-group elements. The difference is that across a row of transition elements, the number of ele's in the outermost principal energy lvl is nearly constant. As another proton is added to the nucleus with each successive element, another ele is added as well, but the ele goes into an n(highest) - 1 orbital. The # of outermost ele's stays constant, and they experience a roughly constant effective nuclear charge, keeping the radius approximately constant. b. As you go down the 1st 2 rows of a column within the transition metals, the elements follow the same general trend in atomic radii and the main-group elements; that is, the radii get larger BC you are adding outermost ele's into higher n lvls.

For transition elements, describe and explain the observed trends in atomic radius as you move: a. across a period in the periodic table b. down a column in the periodic table

(a) Four electron groups give a tetrahedral electron geometry, while three bonding groups and one lone pair give a trigonal pyramidal molecular geometry. (b) Four electron groups give a tetrahedral electron geometry, while two bonding groups and two lone pairs give a bent molecular geometry. (c) Five electron groups give a trigonal bipyramidal electron geometry, while four bonding groups and one pair give a seesaw molecular geometry. (d) Five electron groups give a trigonal bipyramidal electron geometry, while three bonding groups and two lone pairs give a T-shaped molecular geometry. (e) Five electron groups gives a trigonal bipyramidal electron geometry, while two bonding groups and three lone pairs give a linear geometry. (f) Six electron groups give an octahedral electron geometry, while five bonding groups and one lone pair give a square pyramidal molecular geometry. (g) Six electron groups give an octahedral electron geometry, while four bonding groups and two lone pairs give a square planar molecular geometry.

Give the correct electron and molecular geometries that correspond to each set of electron groups around the central atom of a molecule. (a) four electron groups overall; three bonding groups and one lone pair. (b) four electron groups overall; two bonding groups and two lone pairs. (c) five electron groups overall; four bonding groups and one lone pair. (d) five electron groups overall; three bonding groups and two lone pairs. (e) five electron groups overall; two bonding groups and three lone pairs. (f) six electron groups overall; five bonding groups and one lone pair. (g) six electron groups overall; four bonding groups and two lone pairs.

Larger molecules may have two or more interior atoms. When predicting the shapes of these molecules, determine the geometry about each interior atom and use these geometries to determine the entire three-dimensional shape of the molecules.

How do you apply VSEPR theory to predict the shape of a molecule with more than one interior atom?

In a lewis structure, the valence electrons of main-group elements are represented as dots surrounding the symbol for the element. The valence electrons can be determined from the group they are in on the periodic table.

How do you determine how many dots to put around the Lewis symbol of an element?

To determine whether a molecule is polar, do the following: 1. Draw the Lewis structure for the molecule and determine the molecular geometry. 2. Determine whether the molecule contains polar bonds. 3. Determine whether the polar bonds add together to form a net dipole moment.Polarity is important because polar and nonpolar molecules have different properties. Polar molecules interact strongly with other polar molecules but don't interact with nonpolar molecules, and vice versa.

How do you determine whether a molecule is polar? Why is polarity important?

Hybrid orbitals minimize the energy of the molecule by maximizing the orbital overlap in a bond.

How does hybridization of the atomic orbitals in the central atom of a molecule help lower the overall energy of the molecule?

As the ionic radii increases as you move down a group, ions cannot get as close to each other and therefore do not release as much energy when the lattice forms. Thus the lattice energy DECREASES (becomes less negative) as the radius increases.

How does lattice energy relate to atomic radii? To ion charge?

In a solid ionic compound, ions are fixed in place therefore making them nonconductive.In solution, cations/anions dissociate, forming free ions in solution. These ions can move in response to electrical forces, creating an electrical current.

How does the ionic bonding model explain non conductivity of solid ionic compounds and the conductivity of aqueous ionic compounds?

The melting of solid ionic compounds requires enough heat to overcome the electrical forces holding the anions and cations together in a lattice. Thus the melting point is relatively high.

How does the ionic bonding model explain the relatively high melting points of ionic bonds?

-occurs anytime there is a separation of positive and negative charge -quantifies the polarity of a bond -magnitude = qr (charge x distance)

what is a dipole moment?

When an anion is formed from a neutral atom, one valence electron is added for each unit of negative charge. When a cation is formed from a neutral atom, one valence electron is lost for each unit of positive charge. The electron configuration of a main-group monatomic ion can be deduced from the electron configuration of the neutral atom and the charge of the ion. For anions, we simply add the number of electrons required by the magnitude of the charge of the anion. The electron configuration of cations is obtained by subtracting the number of electrons required by the magnitude of the charge.

How is the electron configuration of an anion different from that of the corresponding neutral atom? How is the electron configuration of a cation different?

The number of standard atomic orbitals added together always equals the number of hybrid orbitals formed. The total number of orbitals is conserved.

How is the number of hybrid orbitals related to the number of standard atomic orbitals that are hybridized?

Molecular orbitals can be approximated by a linear combination of atomic orbitals (AOs). The total number of MOs formed from a particular set of AOs will always equal the number of AOs used.

How is the number of molecular orbitals approximated by a linear combination of atomic orbitals related to the number of atomic orbitals used in the approximation?

Nonbinding orbitals are atomic orbitals not involved in a bond that remain localized on the atom.

In molecular orbital theory, what is a nonbonding orbital?

The bond order in a diatomic molecule is the number of electrons in bonding molecular orbitals (MOs) minus the number in antibonding MOs divided by two. The higher the bond order, the stronger the bond. A negative or zero bond order indicates that a bond will not form between the ions.

In molecular orbital theory, what is bond order? Why is it important?

The double bond in Lewis theory is simply two pairs of electrons that are shared between the same two atoms. However, in valence bond theory, we see that the double bond is made up of two different kinds of bonds. The double bond in valence bond theory consists of one sigma bond and one pi bond.Valence bond theory shows that rotation about a double bond is severely restricted. Because of the side-by-side overlap of the p orbitals, the pi bond must essentially break for rotation to occur. The sigma bond consists of end-to-end overlap. Because the overlap is linear, rotation is not restricted.

In the Lewis model, the two bonds in a double bond look identical. However, valence bond theory shows that they are not. Describe a double bond according to valence bond theory. Explain why rotation is restricted about a double bond but not about a single bond.

In valence bonds theory, the interaction energy is usually negative (or stabilizing) when the interacting atomic orbitals contain a total of two electrons that can spin-pair.

In valence bond theory, the interaction energy between the electrons and nucleus of one atom with the electrons and nucleus of another atom is usually negative (stabilizing) when ------.

According to valence bond theory, the shape of the molecule is determined by the geometry of the overlapping orbitals.

In valence bond theory, what determines the geometry of a molecule?

Follow the arrows to get the orbitals in order of increasing energy. The order is 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s

List all orbitals from 1s through 5s according to increasing energy for multi electron atoms

a. The alkali metals (group 1A) have one valence electron and are among the most reactive metals because their outer electron configuration (ns^1) is one electron beyond a noble gas configuration. They react to lose the ns^1 electron, obtaining a noble gas configuration. This is why group 1A metals tend to form +1 cations b. The alkaline earth metals (group 2A) have two valence electrons, have an outer electron configuration of ns^2, and also tend to be reactive metals. They lose their ns^2 electrons to form 2+ cations. c. The halogens (group 7A) have seven valence electrons and have an outer electron configuration of ns^2np^5. They are among the most reactive nonmetals. They are only one electron short of noble gas configuration and tend to react to gain that one electron, forming 1- anions. d. The oxygen family (group 6A) has six valence electrons and has an outer electron configuration of ns^2np^4. They are two electrons short of noble configuration and tend to react to gain those two electrons, forming 2- anions.

List the number of valence electrons for each family in the periodic table, and explain the relationship between the number of valence electrons and the resulting chemistry of the elements in the family. a. Alkali b. Alkaline earth metals c. Halogens d. Oxygen family

The five basic electron geometries: 1. linear, which has two electron groups 2. trigonal planar, which has three electron groups 3. tetrahedral, which has four electron groups 4. trigonal, which has four electron groups 5. octahedral, which has six electron groups An electron group is defined as a lone pair of electrons, a single bond, a multiple bond, or even a single electron.

Name and sketch the five basic electron geometries, and state the number of electron groups corresponding to each. What constitutes an electron group?

(a) A linear electron geometry corresponds to sp hybridization. (b) A trigonal planar electron geometry corresponds to sp² hybridization. (c) A tetrahedral electron geometry corresponds to sp³ hybridization. (d) A trigonal bipyramidal electron geometry corresponds to sp³d hybridization. (e) An octahedral electron geometry corresponds to sp³d² hybridization.

Name the hybridization scheme that corresponds to each electron geometry. (a) linear (b) trigonal planar (c) tetrahedral (d) trigonal bipyramidal (e) octahedral

Lattice energy is associated forming a crystalline lattice of alternating cations and anions from the gaseous ions. Because the cations are positively charged and the anions are negatively charged, there is a lowering of potential- as described by Coulomb's law-when the ions come together to form a lattice energy.That energy is emitted as heat when the lattice forms.

What is lattice energy?

Metals are good conductors of heat and electricity, they are malleable, ductile, often shiny, and tend to lose ele's in chem rxns. As you move to the right across a period in the PT, metallic character decreases. As you move down a column in the PT metallic character increases.

What is metallic character? What are the observed periodic trends in metallic character?

-energy required to break 1 mole of the bond in gas phase -calculate overall enthalpy as sum of enthalpy changes associated with breaking required bonds in reactants and forming required bonds in products

what is bond energy? how can you use average bond energies to calculate enthalpies of reaction?

•based on the observations that properties of elements recur and certain elements have similar properties •explains existence of periodic law is quantum-mechanical theory

The periodic table is a result of the periodic law. What observations led to the periodic law? What theory explains the underly-ing reasons for the periodic law?

As you move to the right across a row in the periodic table, the n level stays the same. However, the nuclear charge increases and the amount of shielding stays about the same because the number of inner electrons stays the same. So the effective nuclear charge experienced by the electrons in the outermost principal energy level increases, resulting is a stronger attraction between the outermost electrons and the nuclease, and therefore, a smaller atomic radii.

Use concepts of effective nuclear charge, shielding, and n value of the valence orbital to explain the trend in atomic radius as you move across a period in the periodic table.

•Degenerate orbitals: orbitals of same energy •in multi electron atom, orbitals in a sub level are degenerate •Hund's rule: when filling degenerate orbitals, electrons fill them singly first with parallel spins •result of an atom's tendency to find the lowest energy state possible

What are degenerate orbitals? According to Hund's rule, how are degenerate orbitals occupied?

•First ionization energy generally increases from left to right across a period. •Except from 2A to 3A and 5A to 6A To ionize N, you must break up a half-full sub level, which costs extra energy. When you ionize O, you get a half-full sub level, which costs less energy.

What are the exceptions to the periodic trends in first ionization energy? Why do they occur?

The three types of bonds are ionic bonds, which occur between metals and nonmetals and are characterized by the transfer of electrons; covalent bonds which occur between nonmetals and are characterized by the sharing of electrons; and metallic bonds which occur between metals and are characterized by electrons being pooled

What are the three basic types of chemical bonds? What happens to electrons in the bonding of atoms in each type?

The electrons that occupy the outer most shell of an atom are called valence electrons. They are important because they determine how an atom will react. By writing an electron configuration, You'll be able to see how many electrons occupy the highest energy level. Electrons in the the outer most orbital, they determine the atom's chemical properties and bonding properties.

What are valence electrons? Why are they important?

•Coulomb's Law: Potential energy (E) of two charged particles depends on their charges (q1 and q2) and separation (r). E=(1/4piE0)(q1q2/r) •PE positive for charges of same sign •PE negative for charges of opposite sign •Magnitude of PE depends inversely on separation between charged particles

What is Coulomb's Law? Explain how potential energy of two charged particles depends on the distance between the charged particles and the magnitude and sign of those charges.

A bonding molecular orbital is lower in energy than the atomic orbitals from which it is formed. There is an increased electron density in the internuclear region.

What is a bonding molecular orbital?

In molecular orbital theory, atoms will bond when the electrons in the atoms can lower their energy by occupying the molecular orbitals of the resultant molecule.

What is a chemical bond according to molecular orbital theory?

According to valence bond theory, a chemical bond results from the overlap of two half-filled orbitals with spin-pairing of the two valence electrons.

What is a chemical bond according to valence bond theory?

An antibonding molecular orbital is higher in energy than the atomic orbitals from which it's formed. There is less electron density in the internuclear region, which results in a node.

What is an antibonding molecular orbital?

•shows the particular orbitals that are occupied by electrons in an atom H= 1s1He=1s2Li= 1s2 2s1

What is an electron configuration. Give an example.

•different way to show electron configuration •symbolizes electron as an arrow in a box that represents the orbital •H, box labeled 1s with an upwards pointing arrow

What is an orbital diagram? Provide an example.

The effective nuclear charge (Zeff) is the average or net charge from the nuclease experienced by the electrons in the outermost level. Shielding is the blocking of nuclear charge from the outermost electrons. The shielding is primarily due to the inner (core) electrons, although here is some interaction and shielding from the electron repulsions of outer electrons with each other.

What is effective nuclear charge? What is shielding?

The EA of an atom or ion is the energy change associated with the gaining of an ele by the atom in the gaseous state. The EA is usually negative BC an atom or ion usually releases energy when it gains an ele. For main-group elements, EA generally becomes more negative as you move to the right across a row in the PT. There is no corresponding trend in the EA going down a column, with the exception of group IA which becomes positive as you go down the column

What is electron affinity? What are the observed periodic trends in electron affinity?

Hybridization is a mathematical procedure in which the standard atomic orbitals are combines to form new atomic orbitals called hybrid orbitals. Hybrid orbitals are still localized on individual atoms, but they have different shapes and energies from those of standard atomic orbitals. They are necessary in valence bond theory because they correspond more closely to the actual distribution of electrons in chemically bonded atoms.

What is hybridization? Why is hybridization necessary in valence bond theory?

The IE of an atom or ion is the energy required to remove an ele from the atom or ion in the gaseous state. The IE is always positive BC removing an ele takes energy. The energy required to remove the first ele is called the first IE. The energy required to remove the second ele is called the 2nd IE. The 2nd IE is always greater than the 1st IE.

What is ionization energy? What is the difference between first ionization energy and second ionization energy?

•The Born-Haber cycle is a hypothetical series of reactions that represents the formation of an ionic compound from its constituent elements. •The reactions are chosen so that the change in enthalpy of each reaction is known except for the last one, which is the lattice energy. •Use Hess's law to add up enthalpy changes of other reactions to determine the lattice energy. •ΔH° (crystal lattice) = lattice energy ▪Don't forget to add together all the ionization energies to get to the desired cation. -hypothetical series of steps that represents formation of an ionic compound from its constituent elements-change in enthalpy for each step is known except for last one, which is lattice energy 1. formation of gaseous sodium from solid sodium 2. formation of chlorine atom from chlorine molecule 3. ionization of gaseous sodium 4. addition of e- to gaseous chlorine 5. formation of crystalline solid from gaseous ions

What is the Born-Haber cycle? List each step in the cycle and show how the cycle is used to calculate lattice energy.

IE generally decreases as you move down a column in the PT BC ele's in the outermost principal energy lvl become farther from the pos charged nucleus and are therefore held less tightly. IE generally increases as you move to the right across a period in the PT BC ele's in the outermost principal energy lvl generally experience a greater effective nuclear charge; therefore, the ele's are closer to the nucleus.

What is the general trend in first ionization energy as you move down a column in the periodic table? As you move across a row?

The electrons in orbitals behave like waves. The bonding molecular orbital arises from the constructive interference between the atomic orbitals and is lower in energy than the atomic orbitals. The antibonding molecular orbital arises from the destructive interference between the atomic orbitals and is higher in energy than the atomic orbitals.

What is the role of wave interference in determining whether a molecular orbital is bonding or antibonding?

When two atomic orbitals are different, the weighting of each orbital in forming a molecular orbital may be different. When a molecular orbital is approximated as a linear combination of atomic orbitals of different energies, the lower-energy atomic orbital makes a greater contribution to the bonding molecular orbital and the higher-energy atomic orbital makes a greater contribution to the antibonding molecular orbital. The shape of the molecular orbital shows a greater electron density at the atom that has the lower atomic orbital energy.

When applying molecular orbital theory to heteronuclear diatomic molecules, the atomic orbitals used may be of different energies. If two atomic orbitals of different energies make two molecular orbitals, how are the energies of the molecular orbitals related to the energies of the atomic orbitals? How is the shape of the resultant molecular orbitals related to the shape of the atomic orbitals?

Cu, Cr

Which of the transition elements in the first transition series have anomalous electron configurations?

-sometimes can write them that are not equivalent -some are better than others -true is represented by average of them, with better weighing more

do resonance structures always contribute equally to the overall structure of a molecule?

-pure covalent/nonpolar: elements w identical electronegativity so they share e- equally -polar covalent: intermediate electronegativity difference between the 2 elements (ex 2 diff nonmetals) -ionic: large electronegativity difference (ex metal and nonmetal) and e- is almost completely transferred

explain the difference between a pure covalent bond, a polar covalent bond, and an ionic bond

-exothermic: when weak bonds break and strong bonds form -endothermic: when strong bonds break and weak bonds form

explain the difference between endothermic reactions and exothermic reactions with respect to the bond energies of the bonds broken and formed

-ratio of a bond's actual dipole moment to the dipole moment it would have if the e- were completely transferred -bonds w e- completely transferred would have 100%, but no bond is 100% ionic -increases as electronegativity difference increases -bonds w more than 50% ionic character are considered ionic bonds

explain what is meant by the percent ionic character of a bond. do any bonds have 100% ionic character?

-when you can write two or more lewis structures for the same molecule -hybrid: weighted average of the structures

what are resonance structures? what is a resonance hybrid?

-3rd row of periodic table and beyond -first or second row never have it

what elements can have expanded octets? what elements should never have expanded octets?