Bonding

• Please read and answer the sheet you have been given.

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Energy and Chemical Bonds• Chemical bonds are the forces that hold atoms

together in a compound.

• Energy is required to overcome these attractive forces and separate the atoms in a compound. Thus, the breaking of a chemical bond is an endothermic process.

• If energy is required to break a bond, then the opposite process of forming a bond must release energy. The formation of a bond is an exothermic process.

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Question

• When a chemical bond is broken, the resulting compound has more potential energy than the substance from which it was formed. Why?– Breaking bonds requires energy, therefore the new

compound will have more energy than at the beginning.

Endothermic

AB + energy → A + B

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Question

• Conversely, when a chemical bond is formed, the resulting compound has less potential energy than the substances from which it was formed. Why?– Forming bonds releases energy, therefore the new

compound will have less energy than at the beginning.

Exothermic

A + B → AB + energy4

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Chemical BondingRECAP

Page 4Complete top half of page.

How can we make Na and Cl happy?

Endothermic or Exothermic?

Exothermic

Chemical BondingRECAP

Page 4Complete bottom half of page.

1.) What is a Chemical Bond

– attractive force between atoms or ions that binds them together as a unit

– bonds form in order to…• decrease potential energy (PE)

• increase stability

MolecularFormula

FormulaUnit

IONIC COVALENT

COCO22NaClNaCl

CHEMICAL FORMULA

COMPOUND

TernaryCompound

BinaryCompound

2 elementsmore than 2

elements

NaNONaNO33NaClNaCl

ION

PolyatomicIon

MonatomicIon

1 atom 2 or more atoms

NONO33--NaNa++

Chemical bonds are formed when valence electrons are:

• transferred from one atom to another (ionic)

• shared between atoms (covalent)• mobile within a metal (metallic)

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Ionic bonds are formed when metals transfer their valence electrons to nonmetals.The oppositely charged ions attract each other to form an ionic bond.

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Sodium has one valence electron and chlorine has seven. Sodium want to lose 1 electron and chlorine needs to gain 1.

Sodium transfers its valence electron to chlorine

Forming an Na+ and a Cl- ion – sodium chloride NaCl

Electron-dot diagrams (Lewis structures) can represent the valence electron arrangement in elements, compounds, and ions.

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atom ion molecular compound

ionic compound

Dots represent valence electrons.Everything else (inner shell electrons and nucleus) is called the Kernel and

is represented by the symbol.

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Phosphorous has 5 valence electrons so we draw 5 dots around the symbol for phosphorous.

Draw the Lewis Dot Structures of the first 18 elements.

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When metals lose electrons to form ions, they lose all their

valence electrons. The Lewis Dot Structure of a metal ion has no dots. The charge indicates how

many electrons were lost.

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Magnesium atom Magnesium ion

When nonmetals gain electrons, they fill up their valence shell with a complete octet (except hydrogen.) The ion is placed in

brackets with the charge outside the brackets.

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A + metal ion is attracted to a – nonmetal ion (opposites attract)

forming an ionic compound. We can use Lewis dot structures to represent

ionic compounds.

23The formula for magnesium fluoride is MgF2

Two major categories of compounds are ionic and

molecular (covalent) compounds.

• Ionic compounds are formed when a metal combines with a nonmetal.

• Ionic compounds have ionic bonds.

• Molecular compounds are formed between two nonmetals.

• Molecular compounds have covalent bonds.

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Comparing the properties compounds with ionic bonds and compounds with covalent bonds.

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Properties of ionic compounds– Solids with high melting

and boiling points (strong attraction between ions)

– Electrolytes: Do not conduct electricity as solids but do when dissolved or molten – ions are charged particles that are free to move

– No individual molecules

Properties of molecular compounds– Low melting and boiling

points (weak attraction between molecules)

– Nonelectrolytes: Do not conduct electricity as solids or when dissolved or molten – no charged particles (ions) to move

– Solids are soft

– Forms molecules

Ionic solids conduct electricity when dissolved or molten.

Molecular solids do not.

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Ionic Solid dissolved in water

Molecular Solid dissolved in water

Solution conducts electricity

Solution doesn’t conduct

electricity

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Nomenclature

“Or How Do We Name Compounds”

Systematic Naming

• Compound is made up of two or more elements

• Name should tell us how many and what type of atoms

• Too many compounds to remember all the names

Anion – Negative ion– Has gained electrons– Non metals form

anions

Cation– Positive ion– Formed by losing

electrons– Metals form cations

Ionic Compounds

• Made of cations and anions• Metals and nonmetals• Electrons lost by the cation are gained by the

anion

Ionic Compounds

Na + Cl

Sodium is cation

1-

ClNa +1+

Chlorine is anion

Charges on Ions

Naming Ions

• Metal ion is written first in both name and formula– It is named directly from element which formed the ion.– Will nearly always be the positive ion or “cation”

– Transition metals can have more than one type of charge– Indicate the charge with roman numerals in parenthesis.

Iron(II) or Iron(III) – Exceptions:

• Silver always +1 • Cadmium and Zinc always +2

Name these

• Na 1+

• Ca 2+

• Al 3+

• Fe 3+

• Fe 2+

• Pb 2+

• Li 1+

• Sodium• Calcium• Aluminum• Iron (III)• Iron (II)• Lead (II)• Lithium

Write Formulas for these

• Potassium ion• Magnesium ion• Copper (II) ion• Chromium (VI) ion • Barium ion• Mercury (II) ion

• K1+

• Mg2+

• Cu2+

• Cr6+

• Ba2+

• Hg2+

Naming Anions

• Anions are always the same.• Change the element ending to -- ide• F1- Fluorine to Fluoride

Name These

• Cl1-

• N3-

• Br 1-

• O2-

• I1-

• Sr2+

• Chloride• Nitride• Bromide• Oxide• Iodide• Strontium

Write These

• Sulfide ion• Iodide ion• Phosphide ion• Strontium ion

• S2-

• I1-

• P3-

• Sr2+

Polyatomic Ions• Tightly bound groups of atoms acting as a

single ion.• Names given in table in book. (pg 123)• Most are anions that contain oxygen. Names

end in –ate (one more O), or –ite (one less O).• SO3

2- = sulfite; SO42- = sulfate

• Exceptions: Ammonium cation NH4+, Cyanide CN-, and hydroxide OH-

Naming Binary Ionic Compounds

• 2 elements involved• Ionic – metal (cation) and a non-metal (anion)• Naming is easy with representative elements

in A groups• NaCl = Na+ Cl- = sodium chloride• MgBr2 = Mg2+Br- = magnesium bromide

Naming Binary Ionic Compounds

• The problem comes with the transition metals.

• Need to figure out their charges• All ionic compounds will have a neutral charge– Same number of + and – charges

• Use the anion to determine the charge on the positive ion.

Naming Binary Ionic Compounds

• Try naming these– KCl– Na3N

– CrN– ScP– PbO– PbO2

– Na2Se

– Potassium chloride– Sodium nitride– Chromium (III) nitride– Scandium (III) phosphide– Lead (II) oxide– Lead (IV) oxide– Sodium selenide

Tertiary Ionic Compounds• Will have polyatomic ions• At least 3 elements• Use blue sheet• Name these ions– NaNO3

– CaSO4

– CuSO3

– (NH4)2O

– LiCN– Fe(OH)3

– (NH4)2CO3

– NiPO4

•Sodium nitrate

•Calcium sulfate

•Copper (II) sulfite

•Ammonium oxide

• Lithium cyanide

• Iron (III) hydroxide

• Ammonium carbonate

• Nickel (III) phosphate

Polyatomic ions are groups of atoms covalently bonded

together that have a negative or positive charge.

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Polyatomic ions are held together by covalent bonds but

form ionic bonds with other ions.

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H N H Cl

H

H

+

-Covalent bonds

Ionic bond

Writing Formulas

• Charges have to add up to zero.• Get charges on pieces from Periodic Table• Cations from element name on table• Anions from table change ending to –ide, or

use name of polyatomic ion• Balance the charges • Put polyatomics in parenthesis

Writing Formulas

• Write formula for calcium chloride– Calcium is Ca2+

– Chloride is Cl1-

– Ca+2Cl-1 would have a +1 charge– Need another Cl1-

– Ca+2Cl2-1 = CaCl2

Writing Formulas• Crisscross method

Ca2+ Cl1- CaCl2No need to write the oneIron (III) sulfide

Calcium chloride

Fe 2 S3

Fe 3+ S2-

Fe2S3

Write Formulas for These

• Lithium sulfide• Tin (II) oxide• Tin (IV) oxide• Magnesium fluoride• Copper (II) sulfate• Iron (III) phosphide• Iron (III) sulfide• Ammonium chloride• Ammonium sulfide

• Li2S• SnO• SnO2

• MgF2

• CuSO4

• FeP• Fe2S3

• (NH4)Cl• (NH4)2S

Things to Look For

• If cations have ( ), the roman numeral is their charge.

• If anions end in –ide they probably are off the periodic table (monoatomic)

• If anion ends in –ate or –ite it is a polyatomic ion

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COVALENT BONDbond formed by the sharing of electrons

Covalent Bond

• Between nonmetallic elements of similar electronegativity.

• Formed by sharing electron pairs• Stable non-ionizing particles, they are not

conductors at any state• Examples; O2, CO2, C2H6, H2O, SiC

Bonds in all the polyatomic ions and diatomics

are all covalent bonds

when electrons are shared equally

NONPOLAR COVALENT BONDS

H2 or Cl2

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

Oxygen AtomOxygen Atom Oxygen AtomOxygen Atom

Oxygen Molecule (O2)

when electrons are shared but shared

unequally

POLAR COVALENT BONDS

H2O

Polar Covalent Bonds: Unevenly matched, but willing to share.

- water is a polar molecule because oxygen is more electronegative than hydrogen, and therefore electrons are pulled closer to oxygen.

when electrons are from one element

only

COORDINATE COVALENT BONDS

NH4+1 or H3O+1

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Homework Page 5

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Molecular Compounds

Writing Names and Formulas

Covalent Bonding / Compounds

• Compounds in which the electronegativity difference is less than 1.7

• Between a nonmetal and nonmetal• Can’t be held together because of opposite

charges• Can’t use charges to figure out how many of

each atom

Covalent Bonding

• Smallest piece of a covalently bonded compound is a molecule

• Electrons are shared between atoms in bond

Water

H2O

Carbon Dioxide

CO2 Ammonia

NH3

In a multiple covalent bond, more than one pair of electrons are shared between two atoms.

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•Diatomic oxygen has a double bond O=O (2 shared pairs) because oxygen needs 2 electrons to fill its valence shell

•Diatomic nitrogen has a triple bond NN (3 shared pairs) because nitrogen needs 3 electrons to fill its valence shell

•Carbon dioxide has two double bonds

Regents Question: 08/02 #17

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Which molecule contains a triple covalent

bond?

(1) H 2

(2) N 2

(3) O 2

(4) Cl 2

Molecular polarity can be determined by the shape of the molecule and the

distribution of charge.• Possible shapes– Linear (X2 HX CO2)

– Bent (H2O)

– Pyramidal (NH3)

– Tetrahedral (CH4 CCl4)

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A polar molecule is called a dipole. It has a positive side and a negative side – uneven charge distribution.

Symmetrical (nonpolar) molecules include CO2 ,

CH4 , and diatomic elements. ..

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Symmetrical molecules are not dipoles.

Asymmetrical (polar) molecules include HCl, NH3 , and H2 O. (5.2l)

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The negative side of the molecule is the side that has the atom with the higher electronegativity.

Differences between ionic and covalent bonding:

Na + Cl-

ClNa ++

Ionic bonding

• electron is “stolen”

• high electronegativity difference

• between metal & nonmetal

• Formation of crystal structure

think proportions of atoms in

formula unit NaCl 1:1

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• Homework Page 10

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Two pages after 11 before 12

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Page 13 for notes

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Molecules are easier to name and work with

• Ionic compounds use charges to determine how many of each.– Have to figure out charges– Have to figure out numbers

• Molecular compound’s name tells you the number of atoms.

Naming

• The second part of all names end with -ide

• Prefixes are used to indicate number of each atom

Prefixes

• 1 mono-• 2 di-• 3 tri-• 4 tetra-• 5 penta-• 6 hexa-• 7 hepta-• 8 octa-

• 9 nona-• 10 deca-

Naming Continued

• To write the name…write two wordsPrefix-name Prefix-name –ide

• One exception is we don’t write mono- if there is only one of the first element.

• No double vowels when writing names– (oa oo)

Name These

• N2O

• NO2

• Cl2O7

• CBr4

• CO2

• BaCl2

• H2O

• Dinitrogen monoxide• Nitrogen dioxide• Dichlorine heptoxide• Carbon tetrabromide• Carbon dioxide• Barium chloride• Dihydrogen monoxide

Write Formulas for These

• Diphosphorous pentoxide• Tetraiodine monoxide• Sulfur hexaflouride• Nitrogen trioxide• Carbon tetrahydride• Phosphorous trifluoride• Aluminum chloride

• P2O5

• I4O

• SF6

• NO3

• CH4

• PFl3

• AlCl3

Page 13 practice

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Page 16 practice

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Page 21 summary

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ionic covalent

valence electrons

Comparison of Bonding Types

sharing of electrons

transfer of electrons

ionsmolecules

EN > 1.7 EN < 1.7

high mp low mp

molten salts conductive

non-conductive

The bonds holding metals together in their crystal lattice

are called metallic bonds.

• All metals have metallic bonds• “Positive ions immersed in a sea of mobile

electrons”– Bonds are between Kernels, leaving the valence

electrons free to move from atom to atom– Mobile electrons give metals the ability to

conduct electricity 88

METALLIC BONDbond found in

metals; holds metal atoms together

very strongly

Metallic Bond

• Formed between atoms of metallic elements• Electron cloud around atoms • Good conductors at all states, lustrous, very

high melting points• Examples; Na, Fe, Al, Au, Co

Metallic Bonds: Mellow dogs with plenty of bones to go around.

Metallic Bond, A Sea of Electrons

Metals Form Alloys

Metals do not combine with metals. They form Alloys which is a solution of a metal in a metal.Examples are steel, brass, bronze and pewter.

Intermolecular Forces

• Weaker than covalent bonds• Weak intermolecular forces – lower boiling point

The stronger the intermolecular forces, the higher the boiling

points and melting points.

• Ionic Solids• Molecules with Hydrogen bonds• Polar molecules• Nonpolar molecules

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Strongest

Weakest

For nonpolar molecules, the greater the mass, the greater the force of attraction.

Dipole-dipole Forces• Polar molecules attract one another when the partial positive

charge on one molecule is near the partial negative charge on the other molecule

• The polar molecules must be in close proximity for the dipole-dipole forces to be significant

• Dipole-dipole forces are characteristically weaker than ion-dipole forces

• Dipole-dipole forces increase with an increase in the polarity of the molecule

Hydrogen Bonds

• Hydrogen bonds are considered to be dipole-dipole type interactions

• Hydrogen bonds vary from about 4 kJ/mol to 25 kJ/mol (so they are still weaker than typical covalent bonds.

• But they are stronger than dipole-dipole and or dispersion forces.

Hydrogen Bonds

Hydrogen Bonds

Van der Waals Forces

• Weak bonds• Liquefy gases• Bonds that combine gas molecules to form liquid• Ex. CO2 – liquid in toy car

- liquid nitrogen• Molecules must be close to each other• Larger atoms have stronger Van-der Waals forces

ion-dipole forces

• Attractive forces between neutral molecules and charged (ionic) compounds

Ion-dipole forces(Ion-Molecule attraction)

•are important in solutions of ionic substances in polar solvents •(e.g. a salt in aqueous solvent)