Bonding

Forces of attraction that hold atoms together making compounds

Chemical symbols Symbols are used to represent elements Either one capital letter, or a capital letter

with a lower case letter Know names and symbols of elements:

– 1 – 30, plus

–Rb, Cs, Sr, Ba, Ag, Au, Cd, Hg, Pt, Ga, Ge, As, Sn, Pb, Se, Br, I, and U

Basic idea...All chemical bonds form

because they impart stability to the atoms involved

lower energy = greater stability

Quick review All types of chemical bonds

involve electrons Valence electrons, the electrons

in the outermost occupied energy level of an atom, are usually the electrons involved in bonding

The representative elements have the same number of valence electrons as their family number in the American system–Example: Mg, column IIA, 2

valence electrons The transition metals all have

two valence electronsns2(n-1)dx

Lewis dot structures are used to represent the valence electrons–each dot represents a valence

electron

–no more than 8 dots total

–no more than 2 dots on a side

–example = Mg: Na.

.

Lewis dot structures of representative elements

The Octet Rule

Atoms will gain, lose, or share electrons in order to achieve an ns2np6 valence configuration

Sizes of atoms

Periodic trend: atomic radii increase moving down a group– Increasing energy level

Periodic trend: atomic radii decrease moving left to right in a period– The charge felt by the valence electrons

becomes larger

• There is a general decrease in atomic radius from left to right, caused by increasing positive charge in the nucleus.

• Valence electrons are not shielded from the increasing nuclear charge because no additional electrons come between the nucleus and the valence electrons.

Sizes of atoms

• For elements that occur as molecules, the atomic radius is half the distance between nuclei of identical atoms.

• For metals, atomic radius is half the distance between adjacent nuclei in a crystal of the element.

Atomic Radius

Atomic Radius

• Atomic radius generally increases as you move down a group.

• The outermost orbital size increases down a group, making the atom larger.

Sizes of ions

Periodic trend: anions are always larger than the atom they were formed from– Electrons repel each other

Periodic trend: cations are always smaller than the atom they were formed from– Fewer electrons to share same positive

nuclear charge

Ionic Radius

• When atoms lose electrons and form positively charged ions, they always become smaller for two reasons:

1.The loss of a valence electron can leave an empty outer orbital resulting in a small radius.

2.Electrostatic repulsion decreases allowing the electrons to be pulled closer to the radius.

• When atoms gain electrons, they can become larger, because the addition of an electron increases electrostatic repulsion.

Ionic Radius

• Both positive and negative ions increase in size moving down a group.

Ionic Radius

• The ionic radii of positive ions generally decrease from left to right.

• The ionic radii of negative ions generally decrease from left to right, beginning with group 15 or 16.

Ionic Radius

Bonding

Forces of attraction that hold atoms together making compounds

Ionization energy The energy needed to remove a

valence electron from an atom A measure of how tightly the

electrons are being held periodic trend

–increases from the bottom up–increases left to right

In general, metals have lower IE than nonmetals–alkali metals are the lowest IE

family

–noble gases are highest IE family

• The energy required to remove the first electron is called the first ionization energy.

Ionization energy

• First ionization energy increases from left to right across a period.

• First ionization energy decreases down a group because atomic size increases and less energy is required to remove an electron farther from the nucleus.

Ionization energy

Ionization energy

• The ionization at which the large increase in energy occurs is related to the number of valence electrons.

Ionization energy

• Removing the second electron requires more energy, and is called the second ionization energy.

• Each successive ionization requires more energy, but it is not a steady increase.

Ionization energy

Electron affinity A measure of how strongly an

element would like to gain an electron

periodic trend – increases from the bottom up

– increases left to right

– ignore the noble gases

Atoms that lose electrons easily have little attraction for additional electrons (and vice versa)– metals have low IE, low EA– Nonmetals have high IE, high EA

Octet rule: when atoms react, they tend to strive to achieve a configuration having 8 valence electrons

This results in some form of bond formation

Periodic trends… As you move from left to right along a

period… Atoms get

…. Smaller Ionization energy goes

…. Up Electron affinity goes

…. Up

Periodic trends… As you move down a group/family Atoms get

…. Larger Ionization energy goes

…. Down Electron affinity goes

…. Down

Check your understanding

The lowest ionization energy is the ____.

A. first

B. second

C. third

D. fourth

The ionic radius of a negative ion becomes larger when:

A. moving up a group

B. moving right to left across period

C. moving down a group

D. the ion loses electrons

Check your understanding

Electron Configuration of Ions

Na 1s22s22p63s1

– will lose one e- to gain ns2np6 configuration

– Na+ 1s22s22p6

S 1s22s22p63s23p4

– will gain 2 e- to gain ns2np6 configuration– S2- 1s22s22p63s23p6

Ionic Bonding Metals lose electrons easily,

nonmetals have a strong attraction for more electrons

metal atoms will lose electrons to nonmetal atoms, causing both to become ions

1. Metals, having lost one or more electrons, become cations (+)

2. Nonmetals, having gained one or more electrons, become anions (-)

3. Opposites attract: the cations and anions are held together electrostaticly

– called “ionic bonds”

In summary...

Ionic bonds are electrostatic attractions between cations and anions formed when electron(s) are transferred from the low IE, EA metal to the high IE, EA nonmetal

Cation (+)

Ionic compound = crystalline solid

Ionic Compounds

High melting points brittle solids nonconducting as solids conduct electricity as liquids

or aqueous

Ionic Compounds As solids, exist in a 3-D repeating pattern

called a crystal “lattice”

the lattice energy is the energy lowering (stability) accomplished by the formation from “free” ions

Also a measure of the energy required to break apart the ionic compound once formed

The greater the lattice energy, the stronger the force of attraction

Bonding

Forces of attraction that hold atoms together making compounds

Ion dissociation Many ionic compounds will

dissolve in water if it results in more stability (lower E) than in the solid ionic compound

the ions “dissociate” from each other

Ex: CaCl2(s) + H2O Ca2+(aq) + 2Cl-(aq)

Ionic Bond Strength

A measure of the attractive force between the ions

smaller atoms = stronger ionic bonds

fewer atom ratio = stronger bond evidence: melting points

Compare the melting points:

KCl : 776oCKI : 723oCsmaller atoms result in

stronger ionic bonds

Compare the melting points:

CaCl2 : 772oCNaCl : 800oC fewer atoms result in

stronger ionic bonds

Bonding

Forces of attraction that hold atoms together making compounds

Covalent Bonding Covalent bonding involves the

sharing of electron pairs usually between two high EA,

high IE nonmetals–both want more e-’s, neither is

willing to lose the e-’s they have

A nonmetal will form as many covalent bonds as necessary to fulfill the octet rule

example: C, with 4 valence e-’s, will form 4 covalent bonds–results in 8 valence e-’s around

the carbon atom at least part of the time

double and triple covalent bonding is a possibility

When does the octet rule

fail?

H, He and Li Helium strives for 2 valence

electrons– 1s2 configuration

Hydrogen will sometimes will share its one electron with another atom, forming a single covalent bond

Lithium will lose its lone valence electron, gaining the 1s2 configuration of He

Be Be will sometimes lose its 2

valence electrons, gaining the Is2 configuration of He

Be will sometimes form 2 covalent bonds, giving it 4 valence electrons–nuclear charge of +4 cannot handle

8 valence electrons

B Boron will often make three

covalent bonds using its three valence electrons–nuclear charge of +5 cannot

handle 8 valence electrons in a stable manner

“organometallic” compounds

Some metals will form covalent compounds with nonmetals–Hg, Ga, Sn, and others

The octet rule is not followed for the metals,but is for nonmetals

Form 2 or more covalent bonds

P, S, Cl, Se, Br, I Elements in the third period and

lower have empty d orbitals there is room for more than 8

valence electrons These elements will at times

make more than 4 covalent bonds

Rules for Drawing structural formulas

1) Determine the central atom, place the other atoms evenly spaced around the outside

2) Count the total number of valence electrons

3) Draw single bonds between the central atoms and each of the outside atoms

4) Complete the octet on the outside atoms by placing electrons in pairs around the outside atoms (lone pairs)

5) Place any remaining electrons on the central atom in pairs

6) If the central atom does not have its minimum number of electrons (usually 8), form double bonds by moving lone pairs off of the outside atoms and drawing them as bonding pairs

Binary Molecular Nomenclature Two nonmetals no charges to balance multiple subscripts possible

–ex: N2O, NO, NO2, N2O4, N2O5

Use prefixes to represent subscripts mono = 1 di = 2 tri = 3 tetra = 4 penta = 5

Hexa = 6 hepta = 7 octa = 8 nona = 9 deca = 10

Rules, continued..

Change second name to end in “ide”

do not use prefixes on the first word if the prefix is “mono”

always use prefixes on the second name

Examples... CO2

carbon = first word subscript = 1, so no prefix oxide = second word subscript = 2, so prefix = di carbon dioxide

Examples... CO carbon = first word subscript = 1, so no prefix oxide = second word subscript = 1, so prefix = mono carbon monoxide

Examples... SF6

1 sulfur, 6 fluorines sulfur hexafluoride P2O5

2 phosphorus, 5 oxygens diphosphorus pentoxide

Examples... Dinitrogen tetroxide di = 2, so two nitrogen’s tetra = 4, so 4 oxygens N2O4

Dihydrogen monoxide H2O! DHMO.org