Chem 151 Chapters 8 & 9

Bonding Pair

A shared pair of electrons is called a bonding pair. We often represent a bonding pair of electrons by a dash to emphasize that it constitutes a chemical bond.

Coulomb's Law

Coulomb's law, which states that the potential energy (E) of two charged particles depends on their charges (q1 and q1) and on there separation (r).

Draw the best Lewis structure for BrO4⁻ and determine the formal charge on bromine. +1 0 +3 +2 -1

0 Explanation: When drawing the Lewis Structure for BrO4⁻ we see that there are 7 different bonds stemming from Br. We take the number of Valence electrons from Br and subtract them from the number of formal electrons that bond together with the oxygen. So you take 7-7 and we get 0.

Match the following. Questions: 1.valence electrons 2. core electrons 3. number of unpaired electrons in Na 4. 0 5. number of unpaired electrons in Ti2 Answers: A. electrons in the outermost shell B. electrons in completed shells C. 1 D. 2 E. number of unpaired electrons in Zn2+

1. Electrons in the outermost Shell 2. Electrons in completed shells 3. 1 4. Number of unpaired electrons in Zn2+ 5. 2

Lone Pair/ Nonbonding Electrons

A pair that is associated with only one atom-- and therefore not involved in bonding-- is called a lone pair. Lone pair electrons are also called nonbonding electrons.

Percent Ionic Character

A quantity called the percent ionic character is the ratio of a bond's actual dipole moment it would have if the electron were completely transferred from one atom to the other, multiplied by 100%. PIC= ((Measured dipole moment of bond)/(Dipole moment if electron were completely transferred)) x 100%

Resonance Structures

A resonance structure is one of two or more Lewis Structures that have the same skeletal formula (the atoms are in the same location), but the difference are the arrangements.

Writing Lewis Structures for Molecular Compounds

1. Write the correct skeletal structure for the molecule. ~First, hydrogen atoms are always terminal. ~Second, put the more electronegative elements in the terminal position and the less electronegative elements (other than hydrogen) in the central position. 2. Calculate the total number of electrons for the Lewis Structure by summing the valence electrons of each atom in the molecule ~If you are writing a Lewis structure for a polyatomic ion, you must consider the charge of the ion when calculating the total number of electrons. 3. Distribute the electrons among the atoms, giving octets (or duets in the case of H) to as many atoms as possible. ~Begin by placing two electrons between every two atoms. These represent the minimum number of bonding electrons. Then distribute the remaining electrons as lone pairs, first to terminal atoms, and then to the central atom, giving octets. 4. If any atoms lack octets, form a double or triple bonds as necessary to give them octets.

Determine the chemical symbols for the neutral elements corresponding to the electronic configurations shown above. Use proper formatting; letter case matters! 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 1s2 2s2 2p3 1s2 2s2 2p6 3s2 3p6

1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 -Kr 1s2 2s2 2p3 - N 1s2 2s2 2p6 3s2 3p6 - Ar

Give the number of core electrons for Cd. 47 48 44 46 45

46 Explanation: Since we have no idea exactly what the valence electrons are since it does not follow the main-group elements, we have to look at the total electron configuration. The electron configuration for Cd is [Kr] 4d10 5s2. We would take the last section, so the 2, and subtract it from the atomic number and this will give you the number of core electrons.

Write the quantum numbers n and ℓ and select all possible values for mℓ for each subshell of the element. 4s2 3d1

4s2 n= 4 l= 0 ml= 0 3d1 n=3 l=2 ml= -2, -1, 0, 1, 2 Explanation: n will always be the coefficient to the subshell as seen above. To find l you follow the rule where s always equals 0, p always equals 1, d always equals 2, and f always equals 3. In order to find ml you start from the negative number of l and count til you get to the positive.

Polar Covalent Bond

A bond between elements that results in unequal sharing of electrons. The bond is said to be polar-- having a positive pole and a negative pole. A polar covalent bond is intermediate in nature between a pure covalent bond and an ionic bond.

Dipole Moment

A dipole moment occurs any time there is a separation of positive and negative charge.

Atomic Radius

A more general term, the atomic radius, refers to a set of average bonding radii determined from measurements on a large number of elements and compounds. ~ Atomic radii peak with each alkali metal. ~As we move down a column (or family) in the periodic table, atomic radius increases. ~As we move to the right across a period (or row) in the periodic table, atomic radius decreases.

Diamagnetic

An atom or ion in which all electrons are paired is not attracted to an external magnetic field- it is instead slightly repelled- and we say that the atom or ion i diamagnetic

Paramagnetic

An atom or ion that contains unpaired electrons is attracted to an external magnetic field, and we say that the atom or ion is paramagnetic.

Nonbonding Atomic Radius/ Van der Waals radius

An atomic radius determined by the solid's density. The van der Waals radius represents the radius of an atom when it is not bonded to another atom.

Electron Configuration

An electron configuration for an atom shows the particular orbitals that electrons occupy for that atom.

Sublevel Energies

E(energy) E(s orbital)< E(p orbital)< E(d orbital)< E(f orbital)

Bonding Atomic Radius/ Covalent Radius

Another way to define the size of an atom, it is defined differently for nonmetals and metals ~Nonmetals: one-half the distance between two of the atoms bonded together ~Metals: one-half the distance between two of the atoms next to each other in a crystal of the metal.

Penetration

As the outer electron undergoes penetration into the region occupied by the inner electrons, it experiences a greater nuclear charge and therefore a lower energy.

Effective Nuclear Charge

As we have seen, we can define the average or net charge experienced by an electron as the effective nuclear charge.

Electron Configurations

As we move to the right across a row, the orbitals fill in the correct order. With each subsequent row, the highest principal quantum number increases by one. Notice that as we move down a column, the number of electrons in the outermost principle energy level (highest n value) remains the same).

Triple Bond

Atoms can also share three electron pairs. The bond is a triple bond. Triple bonds are even shorter and stronger than double bonds.

Core Electrons and Estimating the Effective Nuclear Charge

Core electrons efficiently shield electrons in the outermost principal energy level from nuclear charge, but outermost electrons do not efficiently shield one another from nuclear charge

Hund's Rule

Filling the orbital singly rather than pairing in one orbital right away. Hund's rule states that when filling degenerate orbitals, electrons fill them singly first, with parallel spins.

Choose the bond below that is most polar. H-Br H-F H-Cl C-H H-I

H-F Explanation: In order to be a polar bond there needs to be an unequal sharing of electrons. This coincides with electronegativity, and finding how polar it is by subtracting their electronegativity. Just know that Fluorine has the largest electronegative charge and Hydrogen has one of the smallest, making it the most polar.

Lewis Symbol

In a Lewis symbol, we represent the valence electrons of main-group elements as dots surrounding the abbreviation for the element. ~Each dot represents a valence electron. The dots are placed around the element's symbol with a maximum of two dots per side.

Orbital Diagram

In an orbital diagram the direction of the arrow (pointing up or pointing down) represents the orientation of the electron's spin.

Octet Rule

In either case the bonding atoms obtain stable configuration is usually eight electrons in the outermost shell, this is known as the octet rule.

Chemical Bond

In the Lewis model, a chemical bond is the sharing or transfer of electrons to attain stable electron configurations for the bonding atoms.

Ground State

Lowest energy state

Dmitri Mendeleev

Mendeleev arranged the elements in a table in which mass increases from left to right and elements with similar properties fall in the same columns.

Place the following in order of increasing radius. Br⁻ Na⁺ Rb⁺ Br⁻ < Na⁺ < Rb⁺ Rb⁺ < Na⁺ < Br⁻ Rb⁺ < Br⁻ < Na⁺ Br⁻ < Rb⁺ < Na⁺ Na⁺ < Rb⁺ < Br⁻

Na⁺ < Rb⁺ < Br⁻ Explanation: Any element that is an anion will most of the time have a larger radius than a neutral and/or cation. Atomic Radii increases as you move down a column in the periodic table, hence why the cation of Na is smaller than Rb and why the largest of them all is the anion of Br.

Rank these elements according to first ionization energy. F O C B N Be Ne Li

Ne F N O C Be B Li Explanation: Noble gases are very stable. Of the elements in a given period of the periodic table, the noble gas will require the most energy to ionize because of this stability. Alkali metals tend to form 1 ion. Of the elements in a given period of the periodic table, the alkali metal will require the least energy to ionize because of the stability of the resulting ion. Start by putting the noble gas at the top, the alkali metal at the bottom, and the remaining elements filled in to match the periodic table. Then, the atom with the ns2 electron configuration (in this case Be) is particularly stable and should move up one spot in the ranking. Similarly, the atom with the ns2np3 electron configuration (in this case N) is particularly stable and should move up one spot in the ranking.

Octet

Notice that atoms with eight valence electrons-- which are particularly stable because they have a full outer level-- are easily identified because they have eight dots, an octet.

Using Lewis structures and formal charge, which of the following ions is most stable? OCN⁻ ONC⁻ NOC⁻ NOC⁻ ONC⁻ OCN⁻ None of these ions are stable according to Lewis theory. All of these compounds are equally stable according to Lewis theory.

OCN⁻ Explanation: Not very sure but here are a set of rules that may help... Rules for estimating stability of resonance structures 1. The greater the number of covalent bonds, the greater the stability since more atoms will have complete octets 2. The structure with the least number of formal charges is more stable 3. The structure with the least separation of formal charge is more stable 4. A structure with a negative charge on the more electronegative atom will be more stable 5. Positive charges on the least electronegative atom (most electropositive) is more stable 6. Resonance forms that are equivalent have no difference in stability and contribute equally (eg. benzene)

Periodic Property

One that is predictable based on an elements position within the periodic table.

Rank these elements according to atomic radius. Ca K Be Rb Mg

Rb K Ca Mg Be Explanation: As you move down a group in the periodic table, the elements have outer electrons in successively larger orbitals, giving them a higher probability of being farther from the nucleus. Effective nuclear charge increases as you move from left to right across a period.

Rank the following elements by effective nuclear charge, Zeff, for a valence electron. Rn Ba Pb Bi Po

Rn Po Bi Pb Ba Explanation: This results in a steady increase in Zeff across a period from left to right. So the element that is farthest left, barium, has the lowest Zeff, and the element that is farthest right, radon, has the highest Zeff. The other elements are arranged between barium and radon in the order they go from left to right within the period: barium

Aufbau Principle

Since we know that electrons occupy the lowest energy orbitals available when the atom is in its ground state and that only two electrons (with opposing spins) are allowed in each orbital, we can systematically build up the electron configurations for the elements.

Match the following. Sr-Sr Cs-I Ca-O Se-I C=N metallic bond weakest ionic bond longest covalent bond highest melting point strongest covalent bond

Sr-Sr - metallic bond Cs-I - weakest ionic bond Ca-O - highest melting point Se-I -longest covalent bond C=N -strongest covalent bond

Born-Harber Cycle

The Born-Harber Cycle is a hypothetical series of steps that represents the formation of an ionic compound from its constituent elements.

The Lewis Model

The Lewis Model named after the American chemist G.N. Lewis. In the Lewis model, valence electrons are represented as dots, and we draw Lewis electron-dot structures or Lewis structures to depict molecules. These structures, which are fairly simple to draw, have tremendous predictive power.

The Lewis Symbol of an Anion

The Lewis symbol of an anion is usually written within brackets with the charge in the upper right hand corner outside the brackets.

Electronegativity

The ability of an atom to attract electrons to itself in a chemical bond (which results in polar and ionic bonds) is called electronegativity.

Resonance Hybrid

The actual structure of the molecule is intermediate between the two or more resonance structures and is called a resonance hybrid.

Chemical Properties

The chemical properties of elements are largely determined by the number of valence electrons they contain.

Core Electrons

The core electrons are those in complete principal energy levels and those in complete d and f sublevels.

Trends in Lattice Energies: Ion Size

The magnitude of the lattice energy of the chlorides decreases accordingly, making the formation of the chlorides less exothermic. In other words, we the ionic radii increase as we move down the column, the ions cannot get as close to each other and therefore do not release as much energy when the lattice forms.

Lattice energy

The energy associated with the formation of a crystalline lattice of alternating cations and anions from the gaseous ions.

Formal Charge

The formal charge of an atom in a Lewis Structure is the charge it would have if all bonding electrons were shared equally between the bonded atoms. Formal charge- number of valence electrons- (Number of nonbonding electrons + 1/2 number of bonding electrons)

Trends in Second and Successive Ionization Energies

The ionization energy increases fairly uniformly with each successive removal of an outermost electron, but then takes a large jump with the removal of the first core electrons.

Estimating the Effective Nuclear Charge

To estimate the effective charge experienced nuclear charge we must distinguish between two different types of shielding: 1) the shielding of the outer most electrons by the core electrons 2) the shielding of the outermost electrons by each other.

Shielding

We can think of repulsion of one electron by other electrons as screening, or shielding that electron from the full effects of the nuclear charge.

Double Bond

When two electron pairs are shared between two atoms the resulting bond is a double bond. In general, double bonds are shorter and stronger than single bonds.

Give the ground state electron configuration for Se. [Ar]3d10 4p4 [Ar]4s2 4d10 4p4 [Ar]4s2 3d10 4p4 [Ar]4s2 3d10 [Ar]4s2 3d10 4p6

[Ar]4s2 3d10 4p4 Explanation: the numbers have to equal 16 because that is what the configuration for Se adds up to. Also 3d comes after 4s, not 4d.

Identify the shortest bond. single covalent bond triple covalent bond double covalent bond all of the above bonds are the same length

triple covalent bond Explanation: Make sure you read this correctly! The shortest and strongest bond is a triple bond, and the longest and weakest bond is a single bond. Do not get them mixed up.

Valence Electrons

~An atom's valence electrons are those that are important in chemical bonding. For main-group elements, the valence electrons are those in the outermost principal energy level. ~The elements in a column of the periodic table have similar chemical properties: they have the same number of valence electrons.

The Halogens (Group 7A)

~As expected from periodic trends, the atomic radius and the density increase for each halogen. ~We can see from the melting and boiling points that F and Cl are both gases at room temp. Br is a liquid, and I is a solid. ~All of the halogens are powerful oxidizing agents-- they are readily reduced, gaining electrons from other substances in their reactions.

The Noble Gases (Group 8A)

~As expected from the periodic trends, the atomic radius and the density increases for each successive noble gas, and the ionization energy decreases. ~As their boiling points indicate, all of the noble gases are gases at room temp. and must be cooled to extremely low temperatures before they liquefy. ~The high ionization energies of the noble gases and their completely full outer quantum levels make them exceptionally unreactive.

Summarizing Atomic Radii for Main-Group Elements

~As we move down a column in the periodic table, the principal quantum number (n) of the electrons in the outermost principal energy level increases, resulting in larger orbitals and therefore larger atomic radii. ~As we move to the right across a row in the periodic table, the effective nuclear charge (Zeff) experienced by the electrons in the outermost principal energy level increases, resulting in a stronger attraction between the outermost electrons and the nucleus, and smaller atomic radii.

Atomic Radii and the Translation Elements

~As we move down the first two rows of a column within the translation metals, the elements follow the same general trend in atomic radii as the main-group elements. ~The atomic radii of the transition elements do not follow the same trend as the main-group elements as we move to the right across a row. Instead of decreasing in size, the radii of transition elements stay roughly constant across each row. ~The number of outermost electrons stays constant, and they experience a roughly constant effective nuclear charge, keeping the radius approx. constant.

Metallic Chatacter

~As we move to the right across a row (or period) in the period table, metallic character decreases. ~As we move down a column (or family) in the periodic table, metallic character increases.

Ionic Radii

~Cations are much smaller than their corresponding atoms. ~Anions are much larger than their corresponding atoms. ~We can observe an interesting trend in ionic size by examining the electrons. ~All of these ions have 18 electrons in exactly the same orbitals, but the radius of each ion gets successively smaller. The reason is the progressively greater number of protons.

Types of Chemical Bonds

~Chemical bonds form because the lower the potential energy between the charged particles that compose atoms. ~When a metal bonds with a nonmetal, it transfers one or more electrons to nonmetal. Resulting bond is an ionic bond. ~Therefore when a nonmetal bonds with another nonmetal, neither atom transfers electrons to the other. Instead the two atoms share some electrons. The resulting bond is a covalent bond. ~A third type of bonding, metallic bonding, occurs in metals. Since metals have low ionization energies, they tend to lose electrons easily.

Electron Affinities and Metallic Character

~Electron affinity is a measure of how easily an atom will accept an additional electron and is crucial to chemical bonding because bonding involves the transfer or sharing of electrons. ~EA to become relatively more positive as we move down a column because the electron is entering orbitals with successively higher principal quantum numbers and will therefore be father from the nucleus. This trend applies to the group 1A metals but does not hold for the other columns in the periodic table. ~A more regular trend in electron affinity occurs as we move to the right across a row. ~For the main-group elements, electron affinity generally becomes more negative (more exothermic) as we move to the right across a row in the periodic tabe. The halogens (7A) therefore have the most negative electron affinities.

Main-Group elements and trends in Electronegativity

~Electronegativity generally increases across a period in the periodic table ~Electronegativity generally decreases down a column in the periodic table ~Fluorine is the most electronegative element ~Francium is the least electronegative elements ~Electronegativity is inversely related to atomic size-- the larger the atom, the less ability it has to attract electrons to itself in a chemical bond.

Summarizing Orbital Filling

~Electrons occupy orbitals so as to minimize the energy of the atom; therefore, lower energy orbitals full before higher energy orbitals. Orbitals fill in the following order: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s ~Orbitals can hold no more than two electrons each. When two electrons occupy the same orbital, their spins are opposite. This is another way of expressing the Pauli exclusion principle. ~When orbitals of identical energy are available, electrons first occupy these orbitals singly with parallel spins rather than in pairs. Once the orbitals of equal energy are half full, the electrons start to pair (Hund's rule).

Electrons Configurations and Magnetic Properties of Ions

~For anions, we add the number of electrons indicated by the magnitude of the charge of the anion. ~We determine the electron configuration of cations by subtracting the number of electrons indicated by the magnitude of the charge. ~For main-group cations, we remove the required number of electrons in the reverse order of filling. However for transition metal cations, the trend is different. When writing the electron configuration of a transition metal cations, we remove the electrons in the highest n-value orbitals first, even if this does not correspond to the reverse order of filling. ~In other words, for transition cations, the order in which electrons are removed upon ionization is not the reverse of the filing order.

Conclusion to Coulomb's law

~For like charges, the potential energy (E) is positive and decreases as the particles get father apart. ~For opposite charges, the potential energy is negative and becomes more negative as the particles get closer. ~The magnitude of the interaction between charged particles increases as the charges of the particles increases.

Summarizing Ionization Energy for Main-Group Elements

~Ionization energy generally decreases as we move down a column (or family) in the periodic table because electrons in the outermost principal level are increasingly father away from the positively charged nucleus and are therefore held less tightly. ~Ionization energy generally increases as we move to the right across a row (or period) in the periodic table because electrons in the outermost principal energy level generally experience a greater effective nuclear charge (Zeff).

Summarizing Trends in Lattice Energies

~Lattice energies become less exothermic (less negative) with increasing ionic radius. ~Lattice energies become more exothermic (more negative) with increasing magnitude of ionic charge.

Summarizing EA for Main-Group Elements

~Most groups (columns) of the periodic table do not exhibit any definite trend in electron affinity. Among the group 1A metals, however, electron affinity becomes more positive as we move down the column (adding an electron becomes less exothermic). ~EA generally becomes more negative (adding an electron becomes more exothermic) as we move to the right across a period (row) in the periodic table.

Trends in First Ionization Energy

~Notice the periodic trend in ionization energy, peaking at each noble gas and bottoming at each alkali metal. ~The principal quantum number, n, increases as we move down a column. ~Consequently, electrons in the outermost principal level are farther away from the positively charged nucleus- and are therefore held less tightly- as we move down a column. This results in a lower ionization energy as we move down a column.

The Explanatory Power of the Quantum-Mechanical Model

~Overall energy of the electrons within atoms with eight valence electrons (or two, in the case of Helium) show that they are particularly stable. In other words, when a quantum level is completely full, the overall energy of the electrons that occupy that level is particularly low. ~The noble gases are the most chemically stable and relatively unreactive family in the periodic table. ~Elements with electron configurations close to those of the noble gases are the most reactive because they can attain noble gas electron configurations by losing or gaining a small number of electrons. ~Ex. Alkali metals (group 1A) are among the most reactive metals because their outer electron configuration is one electron beyond a noble gas configuration. ~On the right side of the periodic table, halogens are among the most reactive nonmetals because of their electron configurations. They are only one electron short of a noble gas configuration and tend to react to gain that one electron.

Pauli Exclusion Principle

~Pauli Exclusion Principle: no two electrons in an atom can have the same four quantum numbers. ~The Pauli exclusion principle implies that each orbital can have a maximum of only two electrons, with opposing spins.

Ionization Energy (IE)

~The Ionization Energy of an atom or ion is the energy required to remove an electron from the atom or ion in the gaseous state. Ionization energy is always positive because removing an electrons always takes energy. ~The energy required to remove the first electron is called the first ionization energy (IE1), and so on.

Electronegativity and Bond Polarity

~The arrow with a positive sign on the tail, indicated that the one side of the molecule has a partial positive charge and that the other side of the molecule ( the side the arrow is pointing toward) has a partial negative charge. Similarly the delta plus represents a partial positive and the delta minus represents a partial negative charge. (See page 394 for more detail).

Electron Affinity

~The electron affinity (EA) of an atom or ion is the energy change associated with the gaining of an electron by the atom in the gaseous state. ~ EA is usually-- though not always-- negative because an atom or ion usually releases energy when it gains an electrons.

Bond Polarity

~The greater the electronegativity difference, the more polar the bond. If two atoms with identical electronegativities form a covalent bond, they share the electrons equally, and the bond is purely covalent or nonpolar. ~If there is a large electronegativity differences between the two atoms in a bond, such as normally occurs between a metal and nonmetal, the electron from the metal is almost completely transferred to the nonmetal, and the bond is ionic. ~If there is an intermediate electronegativity difference between the two atoms, such as between two different nonmetals, then the bond is polar covalent.

Summarizing Periodic Table Organization

~The periodic table is divisible into four blocks corresponding to the filling of the four quantum sublevels (s, p, d, f). ~The group number of a main-group element is equal to the number of valence electrons for that element. ~The row number of a main-group element is equal to the highest principal quantum number of that element.

The Transition and Inner Transition Elements

~The principal quantum number of the d orbitals that fill across each row in the transition series is equal to the row number minus one. ~As we move across the f block, the f orbitals fill. For these elements, the principal quantum number of the f orbitals that fill across each row is the row number minus 2. (See pages 348 and 349 for help).

Sublevel Orbital and Electron Relations

~The s sublevel has only one orbital and can therefore hold only 2 electrons ~The p sublevel has three orbitals and can hold 6 electrons ~The d sublevel has five orbitals and can hold 10 electrons ~The f sublevel has seven orbitals and can hold 14 electrons

Concepts for formal charge

~The sum of all formal charges in a neutral molecule must be 0 ~The sum of all formal charges in an ion must be equal to the charge of the ion ~Small (or zero) formal charges on individual atoms are better than large ones ~When formal charge cannot be avoided, negative formal charge should reside on the most electronegative atom.

The Alkali Metals (Group 1A)

~With the exception of potassium, density increases as we move down the column. ~As we move down a column, the increase in mass outpaces the increase in volume caused by greater atomic radius. ~Because of their generally low ionization energies, the alkali metals are excellent reducing agents-- they are readily oxidized, losing electrons to other substances. ~Since ionization energy decreases as we go down the column, the relative reactivities of the alkali metals tend to increase as we move down the column.