1. Bonding.ppt

(1)

(2)

2 • _{Chemical Bond}_{: a mutual electrical}

(3)

3

Ionic Bonding

vs.

Covalent Bonding

• _{Ionic Bonds:}_atoms_{give up}_valence electrons to another atom

• _{Covalent Bonds:}_{atoms use their}

valence electrons to share with other atoms.

(4)

4

Ionic Bond

(5)

(6)

6

Electronegativity and ionic

bonds

• If difference between electronegativities of atoms (p. 151) is more than 1.7, then the bonding has over a 50% ionic

(7)

7

Electronegativity and Covalent

bonding

• If the difference between

electronegativities is 1.7 or less then the bond has ionic characteristics of 50% or less. That means it is covalent and not ionic. The difference between the

(8)

8

Nonpolar covalent bonding

• Covalent bond in which the electrons are shared equally resulting in a balance

distribution of electrical charge. • 0-5% ionic

(9)

Difference in Electronegativity vs.

bonding type

• 0 ≤ nonpolar-covalent ≤ 0.3

• 0.3 < polar-covalent ≤ 1.7

• 1.7 < ionic

(10)

(11)

11

Polar-covalent bonds

• Bonded atoms have an unequal attraction for the shared electrons

(12)

(13)

(14)

14

Classify these sulfur bonds

electronegativity = 2.5

– Electronegativity diff bond type more-neg atom

(15)

15

Classify these sulfur bonds

electronegativity = 2.5

– Electronegativity diff bond type more-neg atom

• H 0.4 polar-cov S • 2.1

• Cs 1.8 ionic S • 0.7

• Cl 0.5 polar-cov Cl • 3.0

(16)

16

Covalent Bonding and

Molecular Compounds

Ch.6.2

(17)

17

• Molecule – neutral group of atoms that are held together by covalent bonds. What’s a covalent bond?

• Molecular compound- chemical compound whose simplest units are molecules

• Chemical formula- indicates the relative numbers of atoms of each kind of

chemical compound by using atomic symbols and numerical subscripts.

• Molecular formula- shows the types and numbers of atoms combined in a single molecule of a molecular compound

(18)

18

Diatomic

molecules

(19)

19

Formation of a covalent bond

(20)

20

• <---> both nuclei repel each other (as do both electron clouds)

• --><-- the nucleus of one atom attracts the electron cloud of the other atom, and vice versa.

(21)

21 • _{Bond length:}_{the average distance}

between two bonded atoms.

• _{Bond energy:} the energy required to break a chemical bond and form neutral isolated atoms.

• Notice the change in potential energy from atoms at equilibrium to atoms separate

(22)

22

• H _____ H______

1s 1s

---

H______ H_______

1s 1s

(23)

23

Octet Rule

• Noble gas atoms exist independently in nature. They have stable electron configuration.

• Noble gasses have 8 valence electrons.

• Other atoms can accomplish this by sharing electrons through covalent bonding.

• Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons,

(24)

24 • Fluorine has 7 valence electrons

• Bonds covalently with another F to form a diatomic molecule F₂

• Each F atoms shares one of its original electrons with the other atom

(25)

25

Electron Dot notation

• Valence electrons of an atom of a

(26)

26

Lewis structure

• Formulas in which atomic symbols

represent nuclei and inner-shell electrons • Dot pairs or dashes between two atomic

symbols represent electron pairs in covalent bonds

(27)

27

Here are some drawings

The dots between the symbols are

shared. The rest around a symbol

(28)

28

The pair of dots that are shared

can be replace by a long dash.

(29)

29

Structural formula

• Same as Lewis structure, but without the unshared electrons

(30)

Lewis Structure Steps

1. Determine type and number of atoms in the molecule

2. Write electron-dot notation for each type of atom

3. Determine the total number of valence electrons in the atoms

4. Arrange the atoms and fill in unshared electrons(Carbon is always in middle,

otherwise the least electronegative atom is central)

(31)

(32)

(33)

33

• Single bond- covalent bond formed by one shared pair of electrons between 2 atoms • Double bond- covalent bond formed from

sharing of 2 pairs of electrons between 2 atoms( written as :: or =)

• Triple bond- you guessed it

• As number increases, bonds get

(34)

(35)

(36)

36

Resonance

• Bonding in molecules or ions that cannot be correctly represented by a single lewis structure

(37)

37

Section 6.3

(38)

(39)

Atom – the smallest unit of matter “indivisible”

(40)

Octet Rule = atoms tend to gain, lose or share electrons so as to have 8 electrons

C would like to

_{N would like to} _{O would like to}

(41)

(42)

(43)

IONIC BOND

bond formed

between

two ions by the

transfer of

(44)

Ionic Compound

• Composed of positive and negative ions that are combined so the numbers of

positive and negative charges are equal • Network of pos and neg ions attracted to

eachother

• Unlike molecular compounds, ionic

compounds not composed of independent, neutral units that can be isolated

(45)

Formula unit

• Is the simplest collection of atoms from

which an ionic compound’s formula can be established.

• F-

and Ca

2+

• CaF₂

(46)

(47)

1). Ionic bond – electron from Na is transferred to Cl, this causes a charge imbalance in each atom. The Na becomes (Na+) and the Cl becomes (Cl-), charged

(48)

• In Ionic crystals, ions minimize their

potential energy by combining in an orderly arrangement known as crystal lattice

• In NaCl- each Na surrounded by six Cl and each Cl surrounded by 6 Na

(49)

(50)

Lattice Energy

• Energy released when one mole of an ionic crystalline compound is formed from

gaseous ions.

• Negative energy values indicate that energy is released when the crystals are formed.

(51)

Ionic vs Molecular compounds

• Strong attraction between positive and negative ions that form ionic bonds are much stronger than attraction between molecules

• Ionic compounds generally have higher melting and boiling points.

(52)

Ionic compounds

• Strong, but brittle

• Slight shift causes like charged ions to line up and thus repel

• In solid ions can not move and so are not conductive

• Molten state and in solution are free to move and carry electric current.

(53)

Polyatomic ions

• Certain atoms bond covalently with each

other to form a group of atoms that has both molecular and ionic characteristics

• A charged group of covalently bonded atoms is a polyatomic ion

• Polyatomic ions combine with ions of

opposite charge to form ionic compounds

(54)

Polyatomic Ions and the Lewis

Structure

• Same steps as in problem 6-4 on page 174 with few exceptions

• If ion is neg charged, add electron(s) to the total number of valence electrons

corresponding to negative charge.

– Ex. 2-ion has 2 additional electrons

• If ion positively charged, subtract from the

total number of valence electrons a number of electrons corresponding to ion’s positive

(55)

6-4

Metallic

Bonding

(56)

COVALENT BOND

bond formed by

(57)

(58)

when electrons

are shared but

shared

unequally

POLAR COVALENT

BONDS

(59)

- water is a polar molecule because oxygen is more

(60)

(61)

62

Polar =

Unequal distribution of charge

(62)

(63)

2. Covalent bonds- Two atoms share one or more pairs of outer-shell electrons.

Oxygen Atom

Oxygen Atom Oxygen AtomOxygen Atom

Oxygen Molecule (O

(64)

when electrons

are shared

equally

NONPOLAR

COVALENT BONDS

(65)

(66)

67 • _{Nonpolar covalent}

bond: bond in which electrons are shared equally between two atoms...very weak.

• _{Polar covalent}

bond: in which one atom attracts the

electrons more than

(67)

METALLIC BOND

bond found in

metals; holds

metal

(68)

• S block metals have open P orbitals in outer level

• D block have 3 open P orbitals and usually many D orbitals of lower level open

• Vacant orbitals overlap, which allows outer

electrons to roam freely throughout entire metal • Electrons are delocalized, they do not belong to

one atom, but move freely about the metal’s network of empty atomic orbitals

(69)

Metallic bonding

• Mobile electrons form a sea of electrons around the metal atoms

• Metal atoms packed together in a crystal lattice

• Metallic bonding is the chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons.

(70)

Metallic Properties

• Freedom of electrons to move promotes conductivity

• Many different orbitals of similar energy levels, so can absorb a wide range of light frequencies

• When electrons return to ground level they emit light in the form of a photon, this is why metals are shiny

(71)

Metallic Properties

• Malleability- ability to be hammered or beaten into thin sheets

• Ductile- ability to be drawn or pulled to form thin wire

• These are possible because metallic

bonding is in direction throughout the solid • One plane of atoms can slide past another

without resistance or breaking bonds

(72)

Heat of Vaporization

• Bond strength vary with nuclear charge of atoms and number of electrons

• When metal vaporized, the atoms in the metal are converted to individual atoms in gaseous state.

• Heat of vaporization used to measure bond strength

(73)

Metallic Bond

• Formed between atoms of metallic elements

• Electron cloud around atoms

• Good conductors at all states, lustrous, very high melting points

(74)

Section 6-5

Molecular

Geometry

(75)

76

Molecular Polarity

• Uneven distribution of molecular charge

(76)

77

VSEPR Theory

• Valence-Shell Electron Pair Repulsion • Repulsion between the sets of

valence-level electrons surrounding an atom

causes these sets to be oriented as far apart as possible

(77)

78

Molecules with no unshared

valence electrons pairs on

central atom

• Three atoms straight line. 180 apart. As far apart as they can get.

• Four atoms 120 deg apart.

• Five atoms 109.5 deg apart. Tetrahedral • Diff surrounding atoms effect angles but

(78)

Table 6-5 p.186

• A

_#

B

_#

E

_#

• A = Central Atom

• B = Atoms around central

Atom

• E = Pairs of unshared

(79)

80

Draw Lewis structure and

predict VSEPR structure of

following

• CBr₄

• AlBr₃

(80)

81 • CBr₄

• AlBr₃

• Be Careful, this one doesn’t make sense because it is ionic. Al and Br have very diff electronegativities. Does it make

sense that metals normally lose electrons? Will Al be neg or pos?

(81)

(82)

83

VSEPR and unshared electron

pairs

• Unshared electrons take up space and effect the position of the bonded atoms

• The actual described shape is only based on the actual position of the atoms

• See figure 6-22 on page 185

(83)

Water molecule

(84)

85

Water molecule

• Still takes form from tetrahedral

• Actual manifested shape is that of a bent line or v

• Angle between legs of tetrahedral

containing atoms is about 105, less than 109.5

– This is because unshared electron pairs repel bonded atoms more strongly

(85)

Water molecule drawn

(86)

Double, triple bonds, and

polyatomic ions

• Double and triple bonds treated the same as single

• Polyatomic ions treated similar to molecules

– Remember to consider all of the electron pairs present in an ion or molecule

(87)

Practice

Predict molecular geometries

• SF₂

• PCl₃

(88)

Hybridization

• VSEPR explain shape but not reveal

relationship between molecule’s geometry and the orbital's occupied by its bonding electrons

• Orbitals of atom become rearranged when atom forms covalent bonds

• The mixing of two or more atomic orbitals of similar energies on the same atom to produce new orbitals of equal energies

(89)

Methane CH

₄

• Tetrahedral geometry

• Carbon has valence electrons with 2 in s and 2 in p

• To achieve 4 equivalent bonds, valence electrons must be the same

• Carbon’s one 2s and three 2p orbitals hybridize to form four new identical orbitals with the same shape

• Four orbital orientations in new sp3 orbital

– 3 indicates that 3 p orientations were used to create new hybridization

• sp3 orbitals all have same energy

– Greater energy than 2s but less than 2p • See figure 6-23 on p. 188

(90)

Hybrid orbitals

• Orbitals of equal energy produced by the combination of 2 or more orbitals on the same atom

• See fig 6-24 p. 188

• Table 6-6 explains how the hybrid orbitals depend on what atoms are involved and gives geometry of diff hybrid orbitals

– Many of these do not follow octet rule

(91)

Intermolecular Forces

• At boiling point, energy is sufficient to

overcome intermolecular forces and pull molecules apart to form gas

• Bp good measure of intermolecular force • Intermolecular forces are much greater for

ionic compounds and metals than for molecular(covalent) bonds

(92)

• Strongest intermolecular forces exist between polar molecules

• Dipole – created by equal, but opposite charges that are separated by a short distance

• Dipole – Dipole Forces:

– Force of attraction between polar molecules – Pos attract neg of adjacent molecules

– BrF higher bp than F-F

(93)

(94)

- water is a polar molecule because oxygen is more

(95)

(96)

(97)

Hydrogen Bonding

• The intermolecular force in which a

hydrogen atom that is bonded to a highly electronegative atom is attracted to an

unshared pair of electrons of an

electronegative atom in a nearby molecule • Hydrogen becomes very positive

• Hydrogen is small and can come into

close contact with unpaired electrons of adjacent molecules

(98)

London dispersion

• Attraction present in polar and nonpolar

• Electrons in constant motion which creates momentary positive and negative poles

• This temporary dipole can induce a dipole in adjacent molecule

• Held together for an instant by weak attraction

• Only attraction present in noble gases and nonpolar molecules

(99)

London dispersion forces

• Since they depend on motion of electrons, force increases as number of electrons

increases

– Increase with increasing atomic or molar mass

• Check bp on table 6-7 p. 190 to verify these two trends

(100)

Draw the Lewis structure and use the

VSEPR theory to predict the molecular

geometry of the following molecules

• CI₄

• As in a C and an I as in eye

(101)

Draw the Lewis structure and use the

VSEPR theory to predict the molecular

geometry of the following molecules

• As in B and CL • BCl₃

(102)

Basic Assessment Questions

Question 1

Determine the ratio of the atoms in the ionic compound formed in each case.

A. aluminum (Al) and fluorine (F)

B. lithium (Li) and oxygen (O)

Topic 7

(103)

Answers

one AL for every three F

two Li for every O

Topic 7

(104)

Write the correct formula for the ionic

compound formed between atoms of each of the following pairs of elements.

A. sodium (Na) and sulfur (S)

Question 2

B. magnesium (Mg) and nitrogen (N)

Topic 7

(105)

Answers

Na₂S

Mg₃N₂

Topic 7

(106)

Additional Assessment Questions

For each of the following atoms, write the formula of the ion the atom is most likely to form and identify that ion as a cation or an anion.

Questions 1

A. bromine (Br), element 35 B. gallium (Ga), element 31

Topic 7

(107)

Answers

A. bromine (Br), element 35 Br– anion

B. gallium (Ga), element 31 Ga3 cation

Topic 7

(108)

• The elements whose natural state is diatomic are: • hydrogen, • bromine, • nitrogen, • oxygen, • fluorine,

• and iodine • chlorine,

Properties of Molecular Substances

Topic 7

Types of Compounds: Basic Concepts

(109)

• Their formulas can be written as:

• H₂,

• Br₂, • N₂,

• O₂, • F₂,

• and I₂, respectively • Cl₂,

Topic 7

(110)

• If two chlorine atoms combine, they share a single pair of electrons, and each atom attains a stable octet configuration.

Topic 7

(111)

• Two oxygen atoms share two pairs of

electrons to form O₂, and two nitrogen atoms share three pairs of electrons to form N₂.

Topic 7

(112)

Allotropes

• The structure of ozone is different from that of diatomic oxygen.

• Although the diatomic form of oxygen, O₂, is most common in our atmosphere, oxygen also exists as O₃—ozone.

• It consists of three atoms of oxygen rather than the two atoms in diatomic oxygen.

Topic 7

(113)

Allotropes

• Molecules of a single element that differ in crystalline or molecular structure are called

allotropes.

Topic 7

(114)

Allotropes

• The properties of allotropes are usually

different even though they contain the same element.

• This is because structure can be more

important than composition in determining properties of molecules.

Topic 7

(115)

Allotropes

• Phosphorus has three common allotropes: white, red, and black.

• All are formed from P₄ molecules that are

joined in different ways, giving each allotrope a unique structure and properties.

Topic 7

(116)

Formulas and Names of Molecular Compounds

• Chemists have devised a naming system for molecular compounds that is based on a much smaller number of rules than there are

compounds.

• Substances are either organic or inorganic. Compounds that contain carbon, with a few exceptions, are classified as organic

compounds.

Topic 7

(117)

• Compounds that do not contain carbon are called inorganic compounds.

• To name these compounds, write out the name of the first nonmetal and follow it by the name of the second nonmetal with its ending changed to -ide.

Formulas and Names of Molecular Compounds

Topic 7

(118)

Naming Organic Compounds

• You have learned that most compounds that contain carbon are organic compounds.

• Organic compounds make up the largest class of molecular compounds known.

• This is because carbon is able to bond to other

carbon atoms in

rings and chains of many sizes.

Topic 7

(119)

Naming Organic Compounds

• The name of even the most complex organic compound is based on the name of a

hydrocarbon, an organic compound that contains only the elements hydrogen and carbon.

Topic 7

(120)

Question 1

Determine the ratio of the atoms in the ionic compound formed in each case.

Topic 7

(121)

Answers

one AL for every three F

two Li for every O

Topic 7

(122)

Write the correct formula for the ionic

compound formed between atoms of each of the following pairs of elements.

Question 2

Topic 7

(123)

Answers

Na₂S

Mg₃N₂

Topic 7

(124)

Types of Compounds: Additional Concepts

Types of Compounds: Additional Concepts Types of Compounds: Additional Concepts Types of Compounds: Additional Concepts

Topic 7

(125)

Forming Chemical Bonds

• Elements tend to react so as to achieve the stable electron configuration of a noble gas, typically an octet of electrons.

• A cation, or positive ion, is formed when an atom loses one or more electrons.

• An anion, or negative ion, is formed when an atom gains one or more electrons.

• The periodic table is useful in predicting the charges of ions typically formed by various atoms.

Topic 7

(126)

Properties of Ionic Compounds and Lattice Energy

• In a solid ionic compound, the positive ions are surrounded by negative ions, and the

negative ions by positive ions.

• The resulting structure is called a crystal

lattice and contains a regular, repeating, three-dimensional arrangement of ions.

Topic 7

(127)

• This arrangement, which involves strong attraction between oppositely charged ions, tends almost always to produce certain

properties, such as high melting and boiling points and brittleness.

Topic 7

(128)

• Ionic compounds are always nonconductors of electricity when solid but good conductors when melted.

• They also act as electrolytes, substances that conduct electric current when dissolved in water.

Topic 7

(129)

• The combination of these conductivity

characteristics is a very good identifier of

ionic compounds, although each characteristic separately is not very reliable.

Topic 7

(130)

• The energy required to separate one mole of the ions of an ionic compound is called lattice energy, which is expressed as a negative

quantity.

• The greater (that is, the more negative) the lattice energy is, the stronger is the force of attraction between the ions.

Topic 7

(131)

• Lattice energy tends to be greater for more-highly-charged ions and for small ions than for ions of lower charge or large size.

Topic 7

(132)

Naming Ionic Compounds

• Certain polyatomic ions, called oxyanions, contain oxygen and another element.

• If two different oxyanions can be formed by an element, the suffix -ate is used for the

oxyanion containing more oxygen atoms, and the suffix -ite for the oxyanion containing

fewer oxygens.

• In the case of the oxyanions of the halogens, the following special rules are used.

Topic 7

(133)

Naming Ionic Compounds

• four oxygens, per + root

+-ate (example: perchlorate, ClO₄–)

• three oxygens, root + -ate

(example: chlorate, ClO₃–)

• two oxygens, root + -ite

(example: chlorite, ClO₂–)

• one oxygen, hypo- + root + -ite

(example: hypochlorite, ClO–)

Topic 7

(134)

For each of the following atoms, write the formula of the ion the atom is most likely to form and identify that ion as a cation or an anion.

Questions 1

A. bromine (Br), element 35 B. gallium (Ga), element 31

Topic 7

(135)

Name the ionic Compounds that have the following formulas.

Question 2

B. KHSO4

A. Mg(NO₃)₂

Topic 7

(136)

Answers

B. KHSO4

A. Mg(NO₃)₂ magnesium nitrate

potassium hydrogen sulfate

Topic 7