chapter 11 chemical bonding forces that hold atoms together

Chapter 11 Chemical Bonding

Forces that hold atoms together

The Nature of Bonding

• There are several major types of bonds. Ionic, covalent and metallic bonds are the three most common types of bonds.

• Covalent bonds – electrons are shared between atoms.

• Ionic bonds – electrons are transferred between atoms, creating cations and anions.

• Metallic bonds – two or more metals bonded together.

The Nature of Covalent Bonding

• There are two different types of covalent bonds, polar covalent and nonpolar covalent.– polar covalent – electrons are not shared

equally between the two bonded atoms. The electrons are pulled toward the more electronegative of the elements.

– nonpolar covalent – electrons are shared equally between the two bonded atoms.

9_12

Li1.0

Na0.9

K0.8

Rb0.8

Cs0.7

Fr0.7

Be1.5

Mg1.2

B2.0

Al1.5

C2.5

Si1.8

N3.0

P2.1

O3.5

S2.5

F4.0

Cl3.0

Ca1.0

Sr1.0

Ba0.9

Ra0.9

Sc1.3

Y1.2

La–Lu1.1–1.2

Ti1.5

Zr1.4

Hf1.3

V1.6

Nb1.6

Ta1.5

Cr1.6

Mo1.8

W1.7

Mn1.5

Tc1.9

Re1.9

Fe1.8

Ru2.2

Os2.2

Co1.8

Rh2.2

Ir2.2

Ni1.8

Pd2.2

Pt2.2

Cu1.9

Ag1.9

Au2.4

Zn1.6

Cd1.7

Hg1.9

Ga1.6

In1.7

Tl1.8

Ge1.8

Sn1.8

Pb1.8

As2.0

Sb1.9

Bi1.9

Se2.4

Te2.1

Po2.0

Br2.8

I2.5

At2.2

IA IIA

IIIB IVB VB VIB VIIB IB IIB

IIIA IVA VA VIA VIIA

VIIIB

H2.1

Ac–No1.1–1.7

Electronegativities

The formation of a bond between two hydrogen atoms.

Source: Andrey K. Geim/High Field Magnet Laboratory/University of Nijmegen

Probability representations of the electron sharing in HF. (a) What the probability map would look like if the two electrons in the H–F bond were shared

equally. (b) The actual situation, where the shared pair spends more time close to the fluorine atom

than to the hydrogen atom.


• Ionic bonds are formed when there is an electronegativity difference (EN) greater than 2.0.

• Polar covalent bonds form when there is a EN between 0.5 and 1.7.

• Nonpolar covalent bonds form when there is a EN between 0 and 0.49.


• If the EN is between 1.7 and 2.0, an ionic bond will form if a metal is one of the elements, and a polar covalent bond will form if only nonmetals or metalloids are present.

The Nature of Covalent Bonding• What type of bond is formed between the

following elements?

• N and O K and F

• Mg and Cl P and F

• C and H

Bond Polarity• Covalent bonding between unlike atoms

results in unequal sharing of the electrons– One end of the bond has larger electron

density than the other

• The result is bond polarity– The end with the larger electron density gets

a partial negative charge– The end that is electron deficient gets a

partial positive charge

H F••

The three possible types of bonds: (a) a covalent bond formed between identical atoms;

(b) a polar covalent bond, with both ionic and covalent components; and

(c) an ionic bond, with no electron sharing.

Dipole Moment• Bond polarity results in an unequal electron

distribution, resulting in areas of partial positive and partial negative charge

• Any molecule that has a center of positive charge and a center of negative charge in different points is said to have a dipole moment (two different poles of charge).

Dipole Moment

• If a molecule has more than one polar covalent bond, the areas of partial negative and positive charge for each bond will partially add to or cancel out each other

• The end result will be a molecule with one center of positive charge and one center of negative charge

• The dipole moment effects the attractive forces between molecules and therefore the physical properties of the substance

(a) The charge distribution in the water molecule. (b) The water molecule behaves as if it had a positive end and a negative end, as indicated by the arrow.

(a) Polar water molecules are strongly attracted to positive ions by their negative ends. (b) They are also strongly attracted to negative ions by their

positive ends.

Polar water molecules are strongly attracted to each other.

Electron Configuration in Ionic Bonding

• Metals tend to lose their valence electrons, leaving a complete octet in their next-lowest energy level.

• Sodium – (1 valence electron) loses 1 electron and becomes Na+1.

• Na ([Ne]3s1) 1e- + Na+1([Ne])• Calcium – (2 valence electrons) loses 2

electrons and becomes Ca+2.• Ca ([Ar]4s2) 2e- + Ca+2([Ar])


• Nonmetals tend to gain or share valence electrons to complete an octet in their highest energy level.

• Oxygen – (6 valence electrons) gains two electrons to become O-2 .

• O ([He]2s22p4) + 2e- O-2 ([He] 2s22p6)• Phosphorus – (5 valence electrons) gains

three electrons to become P-3.• P ([Ne]3s23p3) + 3e- P-3 ([Ne] 3s23p6)

Formation and Properties of Ionic Compounds

• Ionic bonds – forces of attraction that bind cations and anions together.

• Ionic compound – consists of electrically neutral group of ions joined by electrostatic forces.

• Example: Sodium chloride

Formation and Properties of Ionic Compounds

• At room temperature, most ionic compounds are crystalline solids, where ions are arranged in various 3-D patterns.

• Because of the large attractive forces of the ions to each other the compounds become very stable and have high melting points.

The shape of snowflakes results from bonding.

Source: ©Clyde H. Smith/Peter Arnold, Inc.

Sodium Chloride Crystals

The structure of lithium fluoride.

• Scientists have learned that all of the elements within each group behave similarly because they have the same number of valence electrons.

• Valence electrons - # of electrons in the highest occupied energy level of an atom.

• The number of valence electrons is related to the group numbers on the periodic table.


• Group 1 elements = 1 valence electron.• Group 2 elements = 2 valence electrons.• Groups 3-12 elements = 2 valence electrons.• Group 13 elements = 3 valence electrons.• Group 14 elements = 4 valence electrons.• Group 15 elements = 5 valence electrons.• Group 16 elements = 6 valence electrons.• Group 17 elements = 7 valence electrons.• Group 18 elements = 8 valence electrons.


• 1. Multiply the number of valence electrons by the number of moles of each element.

• 2. Add up all the electrons for each of the elements.

• 3. If there is a charge and it is negative, add that number of electrons to the total.

• 4. If there is a charge and it is positive, subtract that number of electrons from the total.

• Total # of electrons should always be an even number!

Determining Valence Electrons for an Ion or a Compound

• Determine the number of valence electrons in each of the following compounds and ions:

• NH4+1

• CH2ClBr

• PO4-3

Determining Valence Electrons Examples

• Valence electrons are the only electrons involved in bonding, and are the only ones written when drawing electron dot structures.

• In forming compounds, atoms tend to achieve the electron configuration of a noble gas, having 8 valence electrons which as known as having a stable octet (octet for 8 valence electrons).


Lewis Symbols of Atoms and Ions• Also known as electron dot symbols• Use symbol of element to represent nucleus and inner

electrons• Use dots around the symbol to represent valence electrons

– put one electron on each side first, then pair• Elements in the same group have the same Lewis symbol

– Because they have the same number of valence electrons

• Cations have Lewis symbols without valence electrons• Anions have Lewis symbols with 8 valence electrons

Li• Be• •B• •C• •N• •O: :F: :Ne:• •

•

• • • •

•• •• •• ••

••

Li• Li+1 :F: [:F:]-1

•

•• ••

••


• Structural formula – chemical formulas that show the arrangement of atoms in molecules and polyatomic ions.

• Octet rule – atoms gain or lose electrons to acquire the stable electron configuration of a noble gas, usually having 8 valence electrons.

Lewis Structures

• You can represent the formation of the covalent bond in H2 as follows:

H .. :H H H+– This uses the Lewis dot symbols for the hydrogen

atom and represents the covalent bond by a pair of dots.

Lewis Structures

• The shared electrons in H2 spend part of the time in the region around each atom.

– In this sense, each atom in H2 has a helium configuration.

:H H

Lewis Structures

• The formation of a bond between H and Cl to give an HCl molecule can be represented in a similar way.

– Thus, hydrogen has two valence electrons about it (as in He) and Cl has eight valence electrons about it (as in Ar).

:H

:

::ClH. . ::

Cl:+

Lewis Structures

• Formulas such as these are referred to as Lewis electron-dot formulas or Lewis structures.

::H Cl:

:– An electron pair is either a bonding pair (shared

between two atoms) or a lone pair (an electron pair that is not shared).

bonding pair

lone pair


• Exceptions to the octet rule: – H needs 2 electrons to be stable– Be needs 4 electrons to be stable– B needs 6 electrons to be stable


• Steps for Drawing Lewis-dot structures1. Determine the number of valence

electrons in the molecule.- When drawing determining valence electrons

for an ion, add electrons if it an anion, and subtract electrons if it is a cation.

2. The first element in the compound will be the central atom. Exception: hydrogen will never be the central atom.


Steps for Drawing Lewis-dot Structures

3. Use one pair of electrons to bond each outer or terminal atom to the central atom.

4. Make all outer or terminal atoms stable using the valence electrons.

5. Put any remaining electrons around the central atom as lone pairs.


• Draw the Lewis structure for:• NH3

• PO43-

• CHFClBr

• PF5-2


• Single covalent bond – a bond in which two atoms share a pair of electrons.

• Double covalent bond – a bond in which two atoms share two pairs of electrons.

• Triple covalent bond – a bond in which two atoms share three pairs of electrons.


• If you have used up all of the valence electrons and you still need two more electrons to make the central atom stable, you must have one double bond.

• If you still need four more electrons to make the central atom stable, you must have either one triple bond or two double bonds.

• Double and triple bonds exist most commonly between C, N, O, and S atoms.


• Draw Lewis structures for:• NOCl

• CO2

• N2

• SiO3-2


• Resonance structures – molecules or ions that can have two or more different Lewis structures. They must contain a double bond to have any resonance structures.

• Resonance structures don’t truly have a single bonds or a double bond, but a hybrid mixture of bonds where the extra bond is spread equally among the other single bonds.

Resonance Structures

• Draw Lewis structures for:• NOCl

• CO2

• N2

• SiO3-2


• Diamagnetic – substance where all of the electrons of the central atom are paired or bonded with other atoms.

• Paramagnetic – substance where all of the electrons of the central atom are not paired or bonded with other atoms.


• Single bonds are longer (length between the atoms) than double and triple bonds.

• Double bonds are longer than triple bonds.

• Single bonds are not as strong as double bonds, and can be broken much easier than double bonds.

• Triple bonds are stronger than double bonds.

Bonding Theory

• The valence-shell electron pair repulsion (VSEPR) model predicts the shapes of molecules and ions by assuming that the valence shell electron pairs are arranged as far from one another as possible.

– To predict the relative positions of atoms around a given atom using the VSEPR model, you first note the arrangement of the electron pairs around that central atom.

Predicting Molecular Geometry

• The following rules and figures will help discern electron pair arrangements.

1. Draw the Lewis structure

2. Determine how many bonding pairs are around the central atom. Count a multiple bond as one pair.

3. Determine how many lone pairs, if any, are around the central atom.

All diatomic molecules have a linear shape.

Arrangement of Electron Pairs About an Atom

3 pairs

Trigonal planar

2 pairs

Linear

4 pairs

Tetrahedral

5 pairs

Trigonal bipyramidal

6 pairs

Octahedral

Molecular Geometry Examples

• NH3

• PO43-

• CHFClBr

• PF5

• SeF6

• NOCl

• CO2

• SF2

• N2

• SiO3-2

Polar Bonds and Molecules

• Nonpolar covalent bond – equal sharing of electrons between two atoms.

• Polar covalent bond – unequal sharing of electrons between two atoms.

• In polar covalent bonds the electrons are pulled closer to the atom with the larger electronegativity value.

• This creates a partial positive and a partial negative pole within the bond.


• Polar bonds can create polar or nonpolar molecules and ions.

• If the centers of partial positive and partial negative charge are in the same location, the molecule or ion is nonpolar.

• If the centers of partial positive and partial negative charge are in different locations, the molecule or ion is polar.


• The easiest way to determine if a molecule or ion is polar or nonpolar is to look at the central atom.

• If the central atom has lone pairs of electrons, the molecule or ion is polar.

• If the central atom does not have any lone pairs of electrons, the molecule or ion is nonpolar.

Examples: Polar or Nonpolar?

• Determine whether each of the following molecules or ions are polar or nonpolar:

• NO2-1

• N2

• CN-1

• CH4

• SO3-2


Attractions Between Molecules

• Molecules are attracted to one another by a variety of forces.

• These intermolecular forces are weaker than ionic or covalent bonds.

• These forces are responsible for whether or not a molecular compound is a solid, liquid, or a gas.


• van der Waals forces – consist of dispersion forces and dipole interactions (dipole-dipole moments).

• Dispersion forces – weakest of all intermolecular forces. They are caused by the motion of electrons. The strength of dispersion forces increases with the increasing number of electrons in a molecule.


• All molecules contain dispersion forces.

• As molar mass and the number of electrons increase, dispersion forces increase.

• Halogens are the most common molecules to have dispersion forces. Fluorine is a gas, Bromine is a liquid and Iodine is a solid.


• Dipole interactions – occur when polar molecules or ions are attracted to one another. This occurs when a partial positive charge and a partial negative charge come close to each other.

• Dipole interactions are very similar to, but much weaker than ionic bonds.


• Hydrogen bonds – force exerted between a hydrogen atom bonded to an F, O, or N atom in one molecule and an unshared pair on another F, O, or N atom in a nearby molecule.

• Hydrogen bonds can have a great effect on the boiling point of a substance.

Intermolecular Forces Examples

chapter 11 chemical bonding forces that hold atoms together

Documents

covalent bonds electrons

bond polarity covalent

covalent components

nonpolar covalent electrons

nonpolar covalent bonds

h slide

ionic bonds electrons

substance slide