Chapter 4 Atoms and Chemical Bonding. Continuous Spectra Roy G. Biv Red Orange Yellow Green Blue Indigo Violet.

Chapter 4

Atoms and Chemical Bonding

Continuous Spectra

Roy G. Biv

RedOrangeYellowGreenBlueIndigoViolet

Line Spectrum

Model of an Atom

• Must explain many things, one of which is the line spectra

• First to come up with a decent model was Bohr• Visualized the electrons in an atom orbiting around

the nucleus like a planet around a sun• Each orbit was associated with a definite energy

(no in-between levels)• Took a “quantum” of energy (fixed amount) to

move from one level to another

Electron Configuration of the AtomBohr (1885-1962)

Electron Energy Levels

Bohr’s Model

• Energy levels each can hold a different (maximum) number of electrons

• Energy level (n) = 1, 2, 3, 4, …• # electrons (2n2)= 2, 8, 18, 32, …

• Example: Na has 11 electrons

– 2e – 8e – 1e (1st) (2nd) (3rd) -> has 17 empty spots in 3rd level

Bohr’s Model

• Worked well for hydrogen• Worked okay for about the next 20 atoms• Didn’t work well at all for anything past the transition

metals• Need new model…

Wave Mechanical Model of the AtomSchrodinger (1881-1961)

Retains Bohr’s energy level concept Distinguishes orbitals at each energy level Orbitals identified as: s, p, d, f

Orbitals in the First Four Energy Levels in the Wave-Mechanical Model of the Atom

Building Atoms with Wave-Mechanical Model

Electrons are added starting at the first level Superscripts are used to indicate the number of

electrons in each orbital

Electron Arrangements of

the First 20 Elements in the Periodic Table

Energy Levels of Oribitals

An Aufbau Diagram

Using the Periodic Table to Determine Electron Configuration

The Shape of an s Orbital

The Shape of a p Orbital

Chemical Properties and the Periodic Table

Chemical properties and electron configuration correlate Alkali metals all have one s electron in their highest

energy level

Li 1s2 2s1

Na 1s22s22p63s1

K 1s22s22p63s23p64s1

Rb 1s22s22p63s23p64s23d104p65s1

Valence ElectronsLewis (1916)

Valence electrons are the electrons in the atom’s highest numbered energy level.

Octet Rule

• Atoms tend to gain or lose electrons to have eight valence electrons• Same as noble gases• He and H are exceptions, get only two valence

electrons

Stable Electron Configurations

• Valence electrons – outermost level with electrons• Core electrons – all other electrons in an atom• Isoelectronic – same number of valence electrons

• Example: O2-, F-, and Ne all have 18 e-’s

Electron-Dot Structures

• Valence electrons represented by dots• Electron-dot symbols

– Examples: Na•, •Mg•, …

Lewis Dot Structures

• The Lewis dot representation (or Lewis dot formulas)• convenient bookkeeping method for valence

electrons• electrons that are transferred or involved in

chemical bonding

Chemical Bonds

• Forces responsible for holding together atoms in molecules and ions in crystals

• Determine shape of molecules• Predict chemical and physical properties of materials• Related to arrangement of electrons in compounds• The electrons involved in bonding are usually those in

the outermost (valence) shell.

Chemical Bonding

• Chemical bonds are classified into two types:• Ionic bonding results from electrostatic attractions

among ions, which are formed by the transfer of one or more electrons from one atom to another.

• Covalent bonding results from sharing one or more electron pairs between two atoms.

Ionic Bonding

• Remember• cations or positive (+) ions

• atoms have lost 1 or more electrons• anions or negative (-) ions

• atoms have gained 1 or more electrons

Ionic Bonding

Atoms lose or gain electrons to form ions Cations are positive ions Anions are negative ions

Ionic compounds are held together by electrostatics- the positive charge of the cation attracting the negative charge of the anion.

Ionic Bonding Continued

Ionic Bonds

• Na+ and Cl–

• Opposite charges attract • Ions organize themselves in orderly manner

• Crystal of NaCl

Structures of Ionic Compounds

• extended three dimensional arrays of oppositely charged ions

• high melting points because coulomb force is strong

Ionic Bonding

• We can also use Lewis formulas to represent the neutral atoms and the ions they form.

Li + F...

.... .

Li+

F[ ]...... ..

Naming Ions

• For cations, simple positive ions• Add the word ion• Examples: Na+ – sodium ion

Al3+ – aluminum ion• For anions, simple negative ions

• Change the usual ending to -ide• Examples: Cl– – chloride

S2– – sulfide

Ionic Compound

Ionic compounds are formed primarily when metals on the left side of the periodic table react with nonmetals on the right side of the periodic table.

Transition metals also form ionic compound Their behavior is less predictable Iron forms Fe2+ or Fe3+

Copper forms Cu+ or Cu2+

Naming Binary Ionic Compounds

• Two components in compound

Common Ions and Their Position in the Periodic Table

Polyatomic Ions

Polyatomic means “many-atom” ion

Mem

orize These

Polyatomic Ions

• “Many-atom” Ions

• Example: NH4+, OH-, CH-

• Covalently bonded groups of atoms, that tend to stay together

Na + O + H O H -- + Na+

Equations

+

+

+

-

-

-

Polyatomic Ions

• Charged groups of atoms that remain together through most chemical reactions

Covalent Bonds

• Bond formed by a shared pair of electrons• Gives atom an octet of electrons

• Shared pair of electrons – bonding pair• Other electrons not involved in bonding –

nonbonding pairs

Covalent Bonding

• covalent bonds formed when atoms share electrons• share 2 electrons - single covalent bond• share 4 electrons - double covalent bond• share 6 electrons - triple covalent bond• attraction is electrostatic in nature

• lower potential energy when bound

Covalent Bonding

extremes in bonding:• pure covalent bonds

• electrons equally shared by the atoms• pure ionic bonds

• electrons are completely lost or gained by one of the atoms

• most compounds fall somewhere between these two extremes

Covalent Bonds

How do two identical atoms bond to form molecules such as H2, N2, or O2?

Neither atom is more likely than the other to transfer an electron

The two atoms have to share electrons

Covalent Bonding

Number of Covalent Bonds/Element

• Follow the electron-dot rules for the following elements

Covalent Bonding

• Lewis dot representation • H molecule formation

+H. H . H H.. or H2

Covalent Bonding

• Lewis dot representation • H molecule formation

• HCl molecule formation

+H. H . H H.. or H2

H Cl H Cl+...

.... ..

..

..

... or HCl

Lewis Dot Structures

• homonuclear diatomic molecules

• hydrogen, H2

• fluorine, F2

• nitrogen, N2

H HorH H..

F F.. .. ....

..

.. ..F F

.. .... ..

.. ..or

N N········ ·· N N·· ··or

Lewis Dot Structures

• heteronuclear diatomic molecules• hydrogen halides

• hydrogen fluoride, HF

• hydrogen chloride, HCl

• hydrogen bromide, HBr

or ··H F··

··H F..

······

or ··H Cl··

··H Cl..

······

or ··H Br··

··H Br..

······

Lewis Dot Structures

• water, H2O

H

H

O··

····

··

Lewis Dot Structures

• ammonia molecule , NH3

H

H

N··

····

·· H

Lewis Dot Structures

• ammonium ion , NH4+

H

H

N··

····

·· H

H +

N - A = S rule

• N = # of electrons needed to be noble gas• 8 for everything except H or He • Only 2 for H or He

• A = # of electrons available in outer shells• equal to group #• 8 for noble gases

N - A = S rule

• for ions• add one e- for each negative charge• subtract one e- for each positive charge

• central atom in a molecule or polyatomic ion is determined by:• atom that requires largest number of

electrons• for two atoms in same group - less

electronegative element is central

Drawing Lewis Dot Formulas

• Example: Write Lewis dot and dash formulas for hydrogen cyanide, HCN.

• N = 2 + 8 + 8 = 18• A = 1 + 4 + 5 = 10• S = 8

Drawing Lewis Dot Formulas

• Example: Write Lewis dot and dash formulas for hydrogen cyanide, HCN.

• N = 2 + 8 + 8 = 18• A = 1 + 4 + 5 = 10• S = 8

H C N·· ·· ···· H C N ··or··

Drawing Lewis Dot Formulas

• Example: Write Lewis dot and dash formulas for the sulfite ion, SO3

2-.

• N = 8 + 3(8) = 32• A = 6 + 3(6) + 2 = 26• S = 6

Drawing Lewis Dot Formulas

• Example: Write Lewis dot and dash formulas for the sulfite ion, SO3

2-.

• N = 8 + 3(8) = 32• A = 6 + 3(6) + 2 = 26• S = 6

O S O

O··

····

····

··

··

··

····

··

····

2-O S

O

O·· ·· ··

······ ··

······

2-or

Naming Covalent Compounds

Use prefixes to indicate the number of each kind of atom

Examples:

Carbon Monoxide

Carbon Dioxide

Trinitrogen Pentoxide

Polar and Nonpolar Covalent Bonds

• nonpolar covalent bonds• electrons are shared equally• symmetrical charge distribution

• must be the same electronegativity to share exactly equally (typically by being the same element)

• H2

• N2

H HorH H..

N N········ ·· N N·· ··or

Polar and Nonpolar Covalent Bonds

• polar covalent bonds• unequally shared electrons• assymmetrical charge distribution• different electronegativities

ElectronegativityPauling (1901-1994)

Electronegativity is the relative tendency of an atom in a molecule to attract a shared pair of electrons in a bond to itself.

The most electronegative element is fluorine and it is given a value of 4.0.

The higher the electronegativity value of an atom, the greater is the ability of an atom of that element to attract electrons to itself.

Electronegativity Values for the Representative Elements

Polar Molecules

When hydrogen and chlorine react to form HCl, a polar molecule is formed.

Continuous Range of Bonding Types

• all bonds have some ionic and some covalent character• HI is about 17% ionic

• greater the electronegativity differences the more polar the bond

Polar Covalent Bonds

• If elements do not have the same electronegativity, they get unequal sharing of electrons

Polar and Nonpolar Covalent Bonds

bondpolar very 1.9 Difference

4.0 2.1 ativitiesElectroneg

F H

1.9

Polar and Nonpolar Covalent Bonds

• Electron density map

of HF• blue areas – low

electron density• red areas – high

electron density• Polar molecules have separation of centers of

negative and positive charge

Polar and Nonpolar Covalent Bonds

bondpolar slightly 0.4 Difference

2.5 2.1 ativitiesElectroneg

I H

0.4

Polar and Nonpolar Covalent Bonds

• Electron density map

of HI• blue areas – low

electron density• red areas – high

electron density• Notice that the charge separation is not as big as for

HF• HI is only slightly polar

Polar Molecules

• Molecule with an overall partial charge• Can have polar bonds and be non-polar

molecule

• O = C = O O

H HNo overall charge - Symmetrical

Overall charge - Asymmetrical

Polar Molecules Continued

Carbon Dioxide, CO2 has polar bonds but because of its symmetrical shape it is a nonpolar molecule.

O=C=O

Intermolecular Forces

• Glue that holds matter together• Melting and boiling points measure the relative

strength • Ionic forces – strongest

• Found in salts• NaCl melts at 800°C

Hydrogen Bonds

Hydrogen Bonding

• Hydrogen must be attached to electronegative atom• N, O, F

• Plays important role in biological systems

Hydrogen Bonding

• Particularly strong dipole-dipole interaction• Occurs between Hydrogen and:

• Oxygen, Nitrogen and Flourine

• Why? High electronegativity of O, N, F

Dipole Moments

• in molecules• some nonpolar molecules have polar bonds

• 2 conditions must hold for a molecule to be polar

units Debye 0.38 units Debye 1.91

I-H F-H

--

Dipole Moments

• There must be at least one polar bond present or one lone pair of electrons.

• The polar bonds, if there are more than one, and lone pairs must be arranged so that their dipole moments do not cancel one another.

Dipole Forces

• Not as strong as ionic forces• Must have polar molecule

• HCl melts at –112°C

Dipole-Dipole Interaction

• Caused by permanent dipoles (partial charges) on the molecule

Cl ClH H

Cl H Cl H- + - +

+ +--

London Forces

Dispersion Forces

• Present in all molecules• Weak momentary attractive forces

• Arise for electrons moving about in molecules and atoms

• Strong in larger molecules• Important in nonpolar compounds

HW Suggestion

• Ch 4:• 3, 4, 6, 9, 11, 13, 18, 19, 20, 23, 24, 27, 29,

33, 34, 35, 37, 39, 40, 43, 45, 46, 49, 53, 55

• Wednesday: Lab! Bring finished prelab writeup, wear closed-toe shoes, hair up, bring goggles, etc…!