Chemical Bonding

Honors Chemistry

Electron Configuration for Ions

• K

• P

• Al

• Se

1[Ar]4s • K+1

• P-3

• Al+3

• Se-2

2 2 6 2 6 11s 2s 2p 3s 3p 4s2 2 6 2 61s 2s 2p 3s 3p

[Ar]

2 2 6 2 31s 2s 2p 3s 3p

2 3[Ne]3s 3p2 2 6 2 61s 2s 2p 3s 3p

2 6[Ne]3s 3p

2 2 6 2 11s 2s 2p 3s 3p

2 1[Ne]3s 3p2 2 61s 2s 2p

[Ne]

2 2 6 2 6 2 10 41s 2s 2p 3s 3p 4s 3d 4p

2 10 4[Ne]4s 3d 4p2 2 6 2 6 2 10 61s 2s 2p 3s 3p 4s 3d 4p

2 10 6[Ne]4s 3d 4p

Periodic Table Trends

• An element’s position & its properties are a result of its electrons

• The outermost electrons, aka valence electrons, have the greatest influence on the properties of the elements.

• Adding an electron to an inner core orbital results in less striking changes in properties than adding an electron to an outer valence orbital (higher energy).

• Shielding Effect: electrons in the lower energy levels (inner core electrons), shield electrons in the outer levels from the full effect of the nuclear charge.

Trends in the Periodic TableA. Atomic Radius

1. The distance from the center of the nucleus to the outermost electron.

2. Atoms get larger going down a group and smaller going across a period.

Ex) Na is larger than Mg

Na is smaller than K

Ga vs. Al

Reasons for trends in Atomic Radius

• Atoms get larger going down a group because each period adds an additional energy level and its more challenging for the nucleus to pull the electrons in closer. (shielding effect)

• Going across a period, atoms become smaller because the nucleus has a greater pull on the electrons.

Note: Atomic volume follows the same trends as atomic radius.

Atomic Radii of the Representative Elements

Atomic Radii vs Atomic Number

Positive Ion Size

• When atoms lose electrons, they become (positive) and get smaller.

• The sizes of cations increases down a group.

• The sizes of cations decreases across a period.

Negative Ionic Size

• When atoms gain electrons, they become (negative) and get larger.

• The sizes of anions increases down a group.

• The sizes of anions decreases across a period.

Relative Sizes of Positive &

Negative Ions

The sodium ion lost an electron, and therefore the positive

protons in the nucleus exert a stronger pull on the remaining negative electrons, shrinking

the orbitals. Thus positive ions are smaller than their atoms.

The chloride ion gained an electron, and therefore the fewer positive protons in the nucleus

exert a weaker pull on the extra negative electrons, increasing the size of the orbitals. Thus

negative ions are larger than their atoms.

• Within an isoelectronic series, radii decrease with increasing atomic number because of increasing nuclear charge.

N-3 O-2 F-1 Na+1 Mg+2 Al+3

How many electrons? Nuclear charge?

> > > > >

Ionic Radius

• Cations & anions decrease in size going across a period

• Cations & anions increase in size going down a group

Electron Attraction in a Bond & Ion Size

Ionization Energy (IE):

1. The energy needed to remove one electron from an atom. (kJ/mole)

2. IE measures how tightly electrons are bound to an atom.– Elements that do not want to lose

their electrons have high ionization energies.

– Elements that easily lose electrons have low ionization energies.

X + energy X+1 + 1 e-

1st Ionization Energy

3. I.E. decreases down a group. 4. I.E. increases across a period

-Note: There will be dips in IE across a period when a sublevel is filled or half-filled.

5. Metals tend to have low IE1.

6. Nonmetals tend to have high IE1.

Ionization Energy of the 1st 20 Elements

1. Energy required to remove electrons beyond the 1st electron.

2. Ionization energies will increase for every electron removed.

X + IE1 X+1 + 1 e-

X+1 + IE2 X+2 + 1 e-

X+2 + IE3 X+3 + 1 e-

Successive Ionization Energies:

Successive Ionization Energies:

kJ/mol IE1 IE2 IE3 IE4 IE5

Na 496 4562 6912 9544 13 353

Mg 738 1451 7733 10 540

13 628

Al 578 1817 2745 11 578

14 831

3s

3s

Electron Configuration

Na: [Ne] 3s1

Mg: [Ne] 3s2

Al: [Ne] 3s23p13p3s

Ionization Energy vs. Atomic Number

Notice the dips across the period… why?

Period 3

Na - [Ne] 3s1 ___ ___ ___ ___Mg - [Ne] 3s2 ___ ___ ___ ___Al - [Ne] 3s23p1 ___ ___ ___ ___Si - [Ne] 3s23p2 ___ ___ ___ ___P - [Ne] 3s23p3 ___ ___ ___ ___S - [Ne] 3s23p4 ___ ___ ___ ___Cl - [Ne] 3s23p5 ___ ___ ___ ___Ar - [Ne] 3s23p6 ___ ___ ___ ___

3s 3p

Electronegativity (EN)

1. Reflects an atoms ability to attract electrons in a chemical bond.• Up to 4.0 for F• Zero for He, Ne, Ar and Kr

2. Metals have low EN.3. Nonmetals have high EN.4. EN decreases down a group.5. EN increases across a period.

Electron Affinity (EA)1. Energy change that occurs when a neutral

gaseous atom gains an electron. Units kJ/mol.1 e- + X X-1 + EA

• Most elements have no affinity for an additional electron and have an EA equal to zero.

He(g) + e- He- EA = 0 kJ/mol

He will not add an electronCl(g) + e- Cl- + 349 kJ/mol EA = -349 kJ/mol

Exothermic!!!

Electron Affinity (EA)

2. Metals have low EA.3. Nonmetals have high EA.4. EA decreases down a group.5. EA increases (becomes more

negative) across a period.• EXCLUDES noble gases• Exceptions: Groups 2 (~0) and 15 (~0

for N and smaller for P to Bi)• Why? Filled s and half filled p

Metallic Character

1. Reflected by those elements that can lose electrons easily.

2. Increases down a group.3. Decreases across a period.4. The most metallic metal is Cesium.5. The most nonmetallic (least

metallic) metal is Aluminum.

Reactivity

• Related to the ability of an element to lose or gain an electron

Brainiac Alkali Metals Alkali Metals Reactivity

1. Reactivity of metals INCREASE down a group.

2. Reactivity of metals DECREASE across a period.

3. Reactivity of nonmetals DECREASE down a group.

4. Reactivity of nonmetals INCREASE across a period.

Introduction to Chemical Bonding

• A chemical bond is an attractive force that holds atoms together in elements or compounds (to function as a unit)– intramolecular (within) vs. intermolecular (between)

• Bond energy is the energy needed to break or form 1 mole of bonds in a gaseous substance (kJ/mol)

• Bonding usually involves only the valence electrons.– In most compounds of the representative elements, the

atoms have an electron configuration that is isoelectronic or psuedoisoelectronic with a noble gas

• The manner in which atoms are bound together in a given substance has a profound effect on its chemical and physical properties.

Octet Rule

• Many chemical compounds form such that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level.– Very stable– There are exceptions: H, He, B (BF3)

• System achieves the lowest possible energy.

Types of Chemical Bonds:

Metallic BondsIonic Bonds

Covalent Bonds

Metallic Bonds

• Simplest crystalline solid – arranged in a very compact and orderly pattern

• Sea of electrons – the valence electrons are mobile around metal cations – Electrons are delocalized

• Attraction of the metal atoms and the surrounding sea of electrons

Metallic Bonds

• Explains metallic properties:–High electrical and thermal conductivity

(flow of electrons)– Luster (metals absorb wide range of -

excites e- and fall back emitting E in form of light results in shiny appearance)

– Ductility & malleability (mobility of e-, metallic bonding is same in all directions throughout solid)

Metallic bonding visual on Holt

Metals vs. Ionic Crystals

• Metallic properties due to sea of electrons

• Ionic compounds are hard but brittle – repulsions result from shift and causes crystal to break

• Chemical bonding that results from an electrostatic attraction between cations and anions to form a neutral compound.

• “salts”• Octet Rule

a) Atoms will transfer electrons (e-) to each other in order to have a full set of valence electrons.

b) When electrons are transferred, ionic bonds are formed.

Ionic bonding

Ionic bonding

Covalent Bonding

• Sharing one or more electron pairs between 2 atoms

Characteristics of Ionic Compounds(hundreds of compounds)

1. All are high melting solids (>400°C).

a) Orderly 3D arrangements (pattern) called crystalline solid or crystal lattice.

b) Simplest arrangement = formula unit

c) High mp reflects strong bonds – large attractive forces are very stable

d) Many are white.e) Colored compounds

usually contain the transition elements (Cu, Cr, Co, Ni, Mn)

2. Solubilitya) Many are soluble in polar solvents, such as

water (aka aqueous solutions)b) Most are insoluble in nonpolar solvents,

such as hexane (C6H14)

Characteristics of Ionic Compounds:

Dissolving Salt Animation

Characteristics of Ionic Compounds:

3. Conductivitya) Solids are non

conductive – ions cannot move freely

b) Molten compounds are conductive – ions move freely (NaCl mp ~800°C)

c) Aqueous solutions are conductive – ions free to move in solution

Formation of Ionic Compounds

• Formation results from a transfer of electrons and the electrostatic attractions of the closely packed, oppositely charged ions.

• Ionic substances are formed when an atom that loses electrons relatively easily reacts with an atoms that has a high affinity for electrons.

• Forms between a metal and a nonmetal. (large EN difference)– Metal: low IE, EN, EA– Nonmetal: high IE, EN, EA– Metal is oxidized (loss of e-) and

nonmetal is reduced (gain of e-)• The ion pair has lower energy than

separated ions.

Electron Configuration

Distribution of electron density• Na: 1s22s22p63s1

– 186 pm

• Cl: 1s22s22p63s23p5

– 99 pm

• Na+1: 1s22s22p6

– 95 pm

• Cl-1: 1s22s22p63s23p6

– 181 pm

Lewis structures or electron-dot structures

Lewis structures examples:

• Sodium and chlorine

• Potassium and phosphorus

1. They are gases, liquids, or solids with low melting points (<300°C)

2. Usually the simplest arrangement is called a molecule (many sizes and shapes)

Characteristics of Covalent Compounds(~11 million compounds)

3. Solubility– Many are insoluble in polar solvents– Many are soluble in nonpolar solvents

4. Conductivity– Liquid and molten compounds are

nonconductors– Aqueous solutions are usually

nonconductors or poor conductors

Characteristics of Covalent Compounds

Formation

1. Formation due to sharing electrons

2. Often forms between nonmetals – metalloid and

nonmetals– small

electronegativity differences

3. Number of covalent bonds likely to form for nonmetals or metalloids depends upon the number unpaired electrons

4. Electron configurationsCl2: 1s22s22p63s23p23p23p1

1s22s22p63s23p23p23p1

1 bond: share 1 pair of e- = single bond

Formation

Electronegativity Trends and Bond Type

• If electronegativity difference is

– Less than .4 = Nonpolar covalent

– Between .4 and 1.9 = Polar Covalent• Can be described a slightly polar (.5 – 1.0)

or moderately polar (1.0 -1.9)

– Greater than 1.9 = Ionic• Can be described as extremely polar

Character of Bonds

ENDifference 0.00 0.65 0.94 1.19 1.43 1.67 1.91 2.19 2.54 3.03

% IonicCharacter 0% 10% 20% 30% 40% 50% 60% 70% 80% 90%

% CovalentCharacter 100% 90% 80% 70% 60% 50% 40% 30% 20% 10%

HOLT

Lewis StructuresGuidelines:1. Select a reasonable skeleton for the

molecule or polyatomic ion. The central atom is the least electronegative atom (excluding H)

2. Calculate the number of shared electrons (S):

S = N – A– N = total number of valence electrons

required (all 8, except H)– A = number of valence electrons available

3. Place shared pairs of electrons in skeleton4. Place lone pairs (for octets)5. NOTE: # of “dots” in the Lewis structure

= A6. For oxyacids, attach the hydrogen to the

oxygen.

Multiple Bonds7 Elements that can multiple bond:

C, N, O, Si, P, S, Se

Examples of Lewis Structures for Covalent Bonds

• Cl2

• ClO4-1

• CO2

Structural formula (line structure) only shows how the molecule or polyatomic ion is bonded – NO “dots” shown

Coordinate Covalent Bonds

• A bond formed when 1 atom provides both electrons (to covalent bond)

• HINTS to determine coordinate covalent bonds:– 1) If there is a negative charge, add those

electrons to the central atom.– 2) If the central atom forms more bonds than

needed, a coordinate covalent bond results.

• Example: Perchlorate (ClO4-1) has 4

coordinate covalent bonds

Equivalent Lewis Structures(A.K.A. Resonance Structures)

• A resonance structure is an alternate way of drawing a Lewis dot structure for a compound.

• For some molecules, there are multiple ways to draw a Lewis dot structure that still satisfy the rules (for instance, having the correct total electron count and satisfying the octet rule on each atom).

Exceptions to the octet rule

• (1) Some stable molecules simply do not have enough electrons to achieve octets around all atoms. This usually occurs in compounds containing Be or B.

• (2) Hypervalence: Elements in the third period and below can accommodate more than an octet of electrons. Although elements such as Si, P, S, Cl, Br, and I obey the octet rule in many cases, under other circumstances they form more bonds than the rule allows.

• (3) Free Radicals: stable molecules which contain an odd number of electrons.

Molecular Shapes

1. Valence shell electron pair repulsion theory (VSEPR Theory): helps to predict the spatial arrangement of atoms in a molecule or polyatomic ion.

a) Introductioni. The central atom is any atom bonded to

more than one other atom.ii. Unshared pairs (lone pair) of electrons and

bonding pairs on the central atom orient themselves to minimize repulsions.

iii. Lone pairs of electrons occupy MORE space than bonding pairs.

VSEPR

b. Counting regions of high electron density around the central atom.

i. Each bonded atom is counted as ONE region of high electron density, whether it is a single, a double, or a triple bond.

ii. Each unshared pair of valence electrons on the central atom is counted as ONE region of high electron density.

Valence Bond (VB) Theory

• Describes HOW bonding occurs• Usually atomic orbitals do not have

the correct energies or orientation to describe where the electrons are when bonded to other atoms

• Hybridization is the mixing of the atomic orbitals to form new hybrid orbitals (s – p – d)

HOLT VSPER

Regions of High Electron Density

• Two regions – LINEAR arrangement• 2 regions e- density = sp• Bond angle = 180°• Example: BeH2

Animation of sp hybridization

Regions of High Electron Density

• Three regions – TRIGONAL PLANAR arrangement

• 3 regions e- density (all bonding) = sp2

• Bond angle = 120°• Example: BF3

Animation of sp2 hybridization

Regions of High Electron Density

• 3 regions e- density = sp2 – 2 bonding and 1 lone pair

• Electronic geometry – trigonal planar• Molecular geometry – BENT or

ANGULAR• Bond angle = 115°• Example: NOCl

Regions of High Electron Density

Regions of High Electron Density

• Four regions – TETRAHEDRAL arrangement

• 4 regions e- density (all bonding) = sp3

• Bond angle = 109.5°• Example: CH4

Regions of High Electron Density

sp3 hybrid orbital

Animation of sp hybridization

Regions of High Electron Density

• 4 regions e- density = sp3

– 3 bonding and 1 lone pairs

• Electronic geometry – tetrahedral• Molecular geometry – TRIGONAL

PYRAMIDAL • Bond angle = 107.3°• Example: NH3

Regions of High Electron Density

Regions of High Electron Density

• 4 regions e- density = sp3

– 2 bonding and 2 lone pairs

• Electronic geometry – tetrahedral• Molecular geometry – BENT or

ANGULAR• Bond angle = 104.5°• Example: H2O

Regions of High Electron Density

Nonpolar Covalent Bonds

• Electron pair is shared equally between the atoms (ΔEN = 0 to ~0.4)– Diatomic molecules (H O F Br I N Cl); allotropes

(S8)

• Electron pair is shared unequally between atoms (ΔEN = ~0.4 to ~1.9) • Results in an electric dipole (2 poles)• Equal but opposite charges that are separated

by a short distance• Separation of charge between 2 covalently

bonded atoms• Examples: HF, HBr, H2O

Polar Covalent Bonds

Character of Bonds

ENDifference 0.00 0.65 0.94 1.19 1.43 1.67 1.91 2.19 2.54 3.03

% IonicCharacter 0% 10% 20% 30% 40% 50% 60% 70% 80% 90%

% CovalentCharacter 100% 90% 80% 70% 60% 50% 40% 30% 20% 10%

HOLT

Polar and Nonpolar Bonds

Polar and Nonpolar Bonds

Ionic, Nonpolar Covalent, Polar Covalent

Molecular Polarity

Consider…a) The presence of at least 1 polar

bond or 1 lone pair of electrons andb) The molecular shape to determine

the overall molecular polarityExamples:HCl, BeCl2, BF3, CH4, NH3, H2O

Molecular Polarity

http://preparatorychemistry.com/Bishop_molecular_polarity.htm