review in chm11-3/2 concepts

REVIEW OF CONCEPTS IN CHM11-‐3/11-‐2

Scien&st and their Important Discovery in Atomic Structure

Scien&st Discovery John Dalton Atom, 3 Fundamental Laws of

Science J.J Thomson Mass/charge raLo of Electron

( -‐ 5 . 6 x 1 0 -‐ 9 g / C ) u s i n g CATHODE RAY TUBE

Robert Millikan Charge of electron (-‐1.602 x 10-‐19 C) using Oil Drop Experiment

Eugene Goldstein Proton through Canal Ray Tube James Chadwick Neutrons using α-‐parLcle at Be

atom Ernest Rutherford N u c l e u s u s i n g

Gold Foil Experiment

ATOMIC STRUCTURE Dalton’s Atomic Theory •  An element is composed of Lny parLcles called atoms – All atoms of the same element have the same chemical properLes

•  In an ordinary chemical reacLon – There is a change in the way atoms are combined with each other

– Atoms are not created or destroyed •  Compounds are formed when two or more atoms of different element combine

Fundamental Laws of Maêr

•  There are three fundamental laws of maêr – Law of conservaLon of mass

•  Maêr is conserved in chemical reacLons

– Law of constant composiLon •  Pure water has the same composiLon everywhere

– Law of mulLple proporLons •  Compare Cr2O3 to CrO3

•  The raLo of Cr:O between the two compounds is a small whole number

Subatomic Particles

Cathode Ray Tube

Millikan (1911) studied electrically-‐charged oil drops.

Charge on each droplet was: n (−1.60 x 10-‐19 C) with n = 1, 2, 3,… n (e-‐ charge)

Modern value = −1.60217653 x 10-‐19 C. = −1 “atomic units”.

These experiments give:

Modern value = 9.1093826 x 10-‐28 g

Note: 1 amu = 1.6605 x 10-‐24 g

= (-‐1.60 x 10-‐19 C)(-‐5.60 x 10-‐9 g/C) = 8.96 x 10-‐28 g

me = charge x mass charge

© 2008 Brooks/Cole

Protons Atoms gain a posiLve charge when e-‐ are lost. Implies a posiLve fundamental parLcle.

Hydrogen ions had the lowest mass. • Hydrogen nuclei assumed to have “unit mass” •  Called protons.

Modern science: mp = 1.67262129 x 10-‐24 g mp ≈ 1800 x me. Charge = -‐1 x (e-‐ charge). = +1.602176462 x 10-‐19 C = +1 atomic units

CANAL RAY TUBE

© 2008 Brooks/Cole

Neutrons Atomic mass > mass of all p+ and e-‐ in an atom. Rutherford proposed a neutral parLcle.

mn ≈ mp (0.1% larger). mn = 1.67492728 x 10-‐24 g.

Present in all atoms (except normal H).

Chadwick (1932) fired α-‐parLcles at Be atoms. Neutral parLcles, neutrons, were ejected:

The Nucleus •  Ernest Rutherford, 1911 •  Bombardment of gold foil with α parLcles (helium atoms minus their electrons – Expected to see the parLcles pass through the foil – Found that some of the alpha parLcles were deflected by the foil

– Led to the discovery of a region of heavy mass at the center of the atom

Nuclear symbolism

•  A is the mass number •  Z is the atomic number •  X is the chemical symbol XAZ

Same element -‐ same number of p+ Atomic number (Z) = number of p+

1 amu = 1.66054 x 10-‐24 g

ParLcle Mass Mass Charge (g) (amu) (atomic units)

e− 9.1093826 x 10-‐28 0.000548579 −1 p+ 1.67262129 x 10-‐24 1.00728 +1 n0 1.67492728 x 10-‐24 1.00866 0

Atomic Numbers & Mass Numbers

Atomic mass unit (amu) = (mass of C atom) that contains 6 p+ and 6

n0.

1 12

For most elements, the percent abundance of its isotopes are constant (everywhere on earth). The periodic table lists an average atomic weight.

Example 1. Naturally occurring chromium consists of four isotopes. It is 4.31% 2450Cr, mass = 49.946 amu, 83.76% 2452Cr, mass = 51.941 amu, 9.55% 2453Cr, mass = 52.941 amu, and 2.38% 2454Cr, mass = 53.939 amu. Calculate the atomic weight of chromium.

Atomic mass = Σ(fracLonal abundance)(isotope mass)

Quantum Theory and Atomic Structure

The Nature of Light

Atomic Spectra

The Wave-‐Par&cle Duality of MaBer and Energy

The Quantum-‐Mechanical Model of the Atom

•  All types (“colors”) have the same velocity (through a vacuum). – c = speed of light = 2.99792458 x 108 ms-‐1 (exact)

•  OscillaLng electric and magneLc fields.

Electric field

MagneLc field

•  Traveling wave § moves through space like the ripples on a pond

ElectromagneLc RadiaLon and Ma^er EM RadiaLon

Wavelength, λ Distance between crests. Length units (m, cm, nm).

distance

λ

Amplitu

de

0

-‐

+

Amplitude Top to bo^om distance.

Frequency, ν = number of crests passing a fixed point per unit Lme. Inverse Lme units (s-‐1).

1 hertz (1 Hz) = 1 s-‐1 νλ = c

ElectromagneLc RadiaLon and Ma^er

γ-‐rays X-‐rays UV IR Microwave Radiowave FM AM Long radio waves

Frequency (Hz) 1024 1022 1020 1018 1016 1014 1012 1010 108 106 104 102 100

10-‐16 10-‐14 10-‐12 10-‐10 10-‐8 10-‐6 10-‐4 10-‐2 100 102 104 106 108 Wavelength (m)

400  450 500 550 600 650 700 Wavelength (nm)

Bacterial Animal Thickness Width Dog cell cell of a CD of a CD

Atom Virus

Visible light is a very small porLon

of the enLre spectrum

E increases from radio waves (low ν, long λ) to gamma rays (high ν, short λ)

ElectromagneLc RadiaLon and Ma^er

Regions of the electromagne&c spectrum.

ElectromagneLc RadiaLon

•  In 1900 Max Planck studied black body radiaLon and realized that to explain the energy spectrum he had to assume that: 1.  energy is quanLzed 2.  light has parLcle character

•  Planck’s equaLon is

22

sJ 10x 6.626 constant s Planck’ h

hc or E h E

34- ⋅==

==λ

ν

Heated solid objects emit visible light •  Intensity and color distribuLon depend on T

Increasing filament T

Planck’s Quantum Theory

24

The Photoelectric Effect

•  Light can strike the surface of some metals causing an electron to be ejected.

An anode (+) a^racts e-‐. Current is measured.

vacuum

Anode (+)

Metal cathode (-‐)

window

“Light” can cause ejecLon of e-‐ from a metal surface.


26


•  What are some pracLcal uses of the photoelectric effect? – Electronic door openers – Light switches for street lights – Exposure meters for cameras

•  Albert Einstein explained the photoelectric effect – ExplanaLon involved light having parLcle-‐like behavior.

– Einstein won the 1921 Nobel Prize in Physics for this work.

The Bohr Model of the Hydrogen Atom •  Heated solid objects emit con(nuous spectra. •  Excited atomic gases emit line spectra. •  Each element has a unique pa^ern.

400  500 600 700 wavelength (nm)

Hydrogen, H

400  500 600 700 wavelength (nm)

Mercury, Hg

The line spectra of several elements.

Flame Color of Some Metal ions METAL ION FLAME COLOR

Na+ Yellow Li+ Red K+ Violet Rb+ Red Cs+ Blue Ca2+ Orange-‐Red Sr2+ Brick red Ba2+ Yellow-‐Green Cu2+ Green

Neils Bohr (1913): •  e-‐ orbit the nucleus. •  e-‐ have fixed E (quanLzed).

•  E-‐levels are idenLfied with integers, n (n = 1…∞).

• Unexcited H atom = e-‐ in the lowest level (n = 1) §  the ground state.

•  Ionized atom (e-‐ removed) has E = 0 (n = ∞).

E = −2.179 x 10-‐18 J n = 1, 2, 3, . . . 1 n2

The Bohr Model of the Hydrogen Atom

Ene

rgy

n ∞ 3 2 1

ultraviolet emission

visible emission

ir emission

400  500 600 700 wavelength (nm)

absorpLon: ΔE > 0, n ↑ emission: ΔE < 0, n ↓

absorpLo

n

Bohr’s model exactly predicts the H-‐atom spectrum.

The Bohr Model of the Hydrogen Atom

De Broglie (1924): all moving objects act as waves:

λ = h mv

λ = wavelength (m) h = Planck’s constant (J s) m = mass (kg) v = velocity (m s-‐1)

Beyond the Bohr Model: Quantum Mechanics

Davidson and Germer (1927) observed e-‐ diffracLon by metal foils.

• Wave-‐like behavior!

Heisenberg Uncertainty Principle It is impossible to know both the exact posiLon and exact momentum of an e-‐.

Why? • Objects only seen by “light” with λ ≤ object size. •  Electrons are very small. Short λ is required.

•  Short λ = high ν = high E. •  EnergeLc collisions alter the speed and direcLon of the e-‐.


Schrödinger equaLon (1926): •  Treats e-‐ as standing waves (not parLcles). • Developed by analogy to classical equaLons for the moLon of a guitar string.

•  Called “wave mechanics” or “quantum mechanics”

§  Explains the structure of all atoms and molecules.

§ Complicated math; important results.


The soluLons are energies and mathemaLcal funcLons (wave funcLons, ψ).

The Schrödinger Equa&on

HΨ = EΨ

d2Ψ dy2

d2Ψ dx2

d2Ψ dz2

+ + 8π2mΘ h2

(E-‐V(x,y,z)Ψ(x,y,z) = 0 +

how ψ changes in space

mass of electron

total quanLzed energy of the atomic system

potenLal energy at x,y,z wave funcLon

An electron density (probability) map plots ψ2 for each point in space. Bigger value = darker shade.


ψ2 = probability of finding an e-‐ at a point in space.

Each ψ describes a different energy level.

Probability maps are hard to draw. •  Boundary surfaces are used

§  contain the e-‐ 90% of the Lme. § Why not 100%? – it would include the enLre universe!

• H-‐atom wavefuncLon (ψ) = an orbital. •  e-‐ do not follow fixed orbits around the nucleus.

The H-‐atom ground-‐state

orbital


38

The Quantum Mechanical Picture of the Atom

Basic Postulates of Quantum Theory 1.  Atoms and molecules can exist only in

certain energy states. In each energy state, the atom or molecule has a definite energy. When an atom or molecule changes its energy state, it must emit or absorb just enough energy to bring it to the new energy state (the quantum condiLon).

39

2.  Atoms or molecules emit or absorb radiaLon (light) as they change their energies. The frequency of the light emi^ed or absorbed is related to the energy change by a simple equaLon.

λν

hch E ==

40

3.  The allowed energy states of atoms and molecules can be described by sets of numbers called quantum numbers.

•  Quantum numbers are the soluLons of the Schrodinger, Heisenberg & Dirac equaLons.

•  Four quantum numbers are necessary to describe energy states of electrons in atoms.

Quantum Numbers Each orbital (ψ) includes three quantum numbers: n, l, and ml

Principal quantum number, n (n = 1, 2, 3, … ∞)

• Most important in determining the orbital energy.

• Defines the orbital size. • Orbitals with equal n are in the same shell.

Quantum Numbers

l 0 1 2 3 4 5 ... Code s p d f g h ...

Azimuthal quantum number, l (l = 0 to n −1)

• Defines the shape of an orbital. • Orbitals with equal l (and equal n) are in the same subshell.

•  Code le^ers idenLfy l

Quantum Numbers MagneLc quantum number, ml (ml = –l to +l ) •  Defines the orientaLon of the orbital. Example List all sets of quantum numbers for an n = 3 e-‐.

Every (n, l, ml) set has a different shape and/or orientaLon.

l = 0, or 1, or 2 if n = 3 and l = 2 (3d), ml is -‐2, -‐1, 0, 1 or 2. if n = 3 and l = 1 (3p), ml is -‐1, 0, or 1. if n = 3 and l = 0 (3s), ml must be 0.

Quantum Numbers Number of Number of Maximum

Electron Subshell Orbitals Electrons Electrons Shell type Available Possible for nth Shell (n) (=2l + 1) in Subshell (=2n2) 1 s 1 2 2 2 s 1 2

p 3 6 8 3 s 1 2

p 3 6 d 5 10 18

4 s 1 2 p 3 6 d 5 10 f 7 14 32

5 s 1 2 p 3 6 d 5 10 f 7 14 g* 9 18 50

Electron Spin Experiments showed a 4th quantum no. was needed

•  +½ or −½ only. spin quantum number, ms

View an e-‐ as a spinning sphere. Spinning charges act as magnets.

•  Pauli: every e-‐ in an atom must have a unique set of (n, l, ml, ms). § Maximum of 2e-‐ per orbital (opposite spins). § Pauli exclusion principle.

s Orbitals l = 0 orbital: Every shell (n level) has one s orbital.

Distance from nucleus, r (pm)

Prob

ability of fi

nding e-‐ at

distance r from

nucleus

Spherical. Larger n value = larger sphere

1s 2s 3s

p Orbitals

Three p orbitals (l = 1): px, py and pz Related to ml = -‐1, 0, +1.

Five d orbitals (l = 2): 3dxz 3dxy 3dyz 3dx -‐ y 3dz 2 2 2

d Orbitals

Periodic Table

•  Au�au Building up Principle •  Pauli’s Exclusion Principle •  Hund’s Rule of Maximum MulLplicity

– MulLplicity – number of unpaired electron plus 1 or the number of possible energy levels that depend on the orientaLon of the net magneLc moment in a magneLc field

Increasing (n + l), then increasing n

1s

2s

3s

4s

5s

6s

7s

8s

2p

3p

4p

5p

6p

7p

3d

4d

5d

6d

5f

4f

n value

8

7

6

5

4

3

2

1

l value

0 1 2 3

n + l = 1

n + l = 2 n + l = 3

n + l = 4 n + l = 5

n + l = 6 n + l = 7

n + l = 8

Atom Electron ConfiguraLons

Main group s block

2 s

4 s

5 s

6 s

7 s

3 s

1s

5 f

4 f

6d

4d

3d

5d

6d

5d

4d

3d 4p

5p

6p

7p

3p

2p

1s

Lanthanides and acLnides f block

TransiLon elements d block

Main group p block

Block idenLLes show where successive e-‐ add. Note: d “steps down”, f “steps down” again.


H

He

Li

Be

B

C

N

O

F

Ne

1s 2s 2p Electron configuraLons Expanded Condensed

1s1 1s1

1s2 1s2

1s22s1 1s22s1

1s22s2 1s22s2

1s22s22p1 1s22s22p1

1s22s22p12p1 1s22s22p2

1s22s22p12p12p1 1s22s22p3

1s22s22p22p12p1 1s22s22p4

1s22s22p22p22p1 1s22s22p5

1s22s22p22p22p2 1s22s22p6

1s

2s 2p

3s 3p

Energy


§  The magnets cancel. §  The atom is diamagneLc

•  pushed weakly away from magneLc fields.

Spinning e-‐ = Lny magnet. If all e-‐ are paired:

With unpaired e-‐: §  Unpaired spins point in the same direcLon (Hund’s rule). §  Magnets add. §  The atom is paramagneLc.

•  a^racted to magneLc fields.

If individual atom-‐magnets line up in a bulk sample §  A ferromagnet -‐ a permanent magnet. Ex. Fe, Ni, Co

ParamagneLsm & Unpaired Electrons

Paramagnet Ferromagnet ParamagneLsm & Unpaired Electrons

•  As a result of penetraLon and shielding, the order of energies in many-‐electron atoms is typically:

ns < np < nd < nf

•  EsLmate of atomic size – ½(homonuclear bond length) – Cl = 100 pm (Cl2 bond = 200 pm) – H = 37 pm (H2 bond =74 pm)

Cl

Cl

200 pm

100 pm •  Radii are addiLve.

§  HCl has a (37 + 100) = 137 pm bond

Periodic Trends: Atomic Radii

Atoms grow down a group. •  Larger shell (larger n) added in each new row.

Atoms shrink across a period •  e-‐ add to the same shell. Increase size? No,

Big Jump from noble gas to alkali metal • A new shell (with larger n) is added.

•  p+ add to the nucleus. •  Larger charge pulls all e-‐ in, shrinking the atom.

Periodic Trends: Atomic Radii

A caLon is smaller than its neutral atom.

Group 1A Group 2A Group 3A Li Li+ Be Be2+ B B3+

152 90 112 59 85 25

Na Na+ Mg Mg2+ Al Al3+

186 116 160 86 143 68

K K+ Ca Ca2+ Ga Ga3+ 227 152 197 114 135 76

Rb Rb+ Sr Sr2+ In In3+ 248 166 215 132 167 94

Main block: outer shell completely removed. e-‐/e-‐ repulsion reduced (fewer e-‐ ).

Periodic Trends: Ionic Radii

An anion is larger than its neutral atom.

Group 6A Group 7A

O O2-‐ F F-‐ 73 126 72 119

S S2-‐ Cl Cl-‐ 103 170 100 167

Se Se2-‐ Br Br-‐ 119 184 114 182

Te Te2-‐ I I-‐ 142 207 133 206

More e-‐/e-‐ repulsion (more e-‐). The shell swells.

Periodic Trends: Ionic Radii

Periodic Trends: Ionic Radii Isoelectronic Ions O2-‐ F-‐ Na+ Mg2+

Ionic radius (pm) 126 119 116 86 Number of protons 8 9 11 12 Number of electrons 10 10 10 10

Increasing nuclear charge

decreasing size

Polarizability

•  Ability of an atom to be distorted by an electric field (such as that of the neighboring ion

•  Large, heavy atoms and ions tends to be highly polarizable.

IonizaLon energy (IE) E to remove an e-‐ from a gas-‐phase atom.

Mg(g) Mg+(g) + e-‐ ΔE = IonizaLon Energy

Periodic Trends: IonizaLon Energies

Also called IE1

Across a period: IE ↑. Smaller atom = more Lghtly held e-‐ Down a group: IE ↓. Larger atom = less Lghtly held e-‐


Second ionizaLon energy (IE2) E required to remove a 2nd e-‐ from A+(g).

•  IE2 > IE1 • Mg+ holds the 2nd e-‐ more Lghtly. • Huge increase if e-‐ removal breaks a complete shell (the core).

Mg(g) Mg+ (g) + e-‐ ΔE = IE1 Mg+(g) Mg2+(g) + e-‐ ΔE = IE2


Table 7.8 IonizaLon Energies Required to Remove Successive Electrons

IonizaLon Energy Li Be B C N O F Ne (MJ/mol) 1s22s1 1s22s2 1s22s22p1 1s22s22p2 1s22s22p3 1s22s22p4 1s22s22p5 1s22s22p6

IE1 0.52 0.90 0.80 1.09 1.40 1.31 1.68 2.08 IE2 7.30 1.76 2.43 2.35 2.86 3.39 3.37 3.95 IE3 11.81 14.85 3.66 4.62 4.58 5.30 6.05 6.12 IE4 21.01 25.02 6.22 7.48 7.47 8.41 9.37 IE5 32.82 37.83 9.44 10.98 11.02 12.18 IE6 47.28 53.27 13.33 15.16 15.24 IE7 64.37 71.33 17.87 20.00 IE8 Core electrons 84.08 92.04 23.07 IE9 106.43 115.38 IE10 131.43


Electron Affinity (EA) E released when an e-‐ adds to a gas-‐phase atom.

F(g) + e-‐ F-‐(g) ΔE = Electron Affinity

• Usually exothermic (EA is negaLve)

•  EA increases from le� to right.

• Metals have low EA; nonmetals have high EA.

•  Some tables list posiLve numbers.

§ a sign-‐convenLon choice.

Periodic Trends: Electron AffiniLes

Table 7.9 Electron AffiniLes (kJ/mol) 1A 2A 3A 4A 5A 6A 7A 8A (1) (2) (13) (14) (15) (16) (17) (18)

H -‐73

Li -‐60

Na -‐53 K -‐48 Rb -‐47

Ne >0 Ar >0

Kr >0 Xe >0

He >0

Be >0 Mg >0 Ca -‐2 Sr -‐5

B -‐27 Al -‐43 Ga -‐30

In -‐30

C -‐122 Si -‐134 Ge -‐119 Sn -‐107

N >0 P -‐72 As -‐78 Sb -‐103

O -‐141

S -‐200 Se -‐195 Te -‐190

F -‐328

Cl -‐349 Br -‐325 I -‐295

Periodic Trends: Electron AffiniLes

Trends in three atomic proper&es.

ATOMS

Dalton's Atomic Theory

Subatomic Particles

Quantum Theory

Dualistic Nature of

MatterSchrodinger

Equation

Quantum Numbers

Wave Function

Atomic Orbitals

Protons

Neutrons

Electrons Periodic Table

Aufbau Principle

Pauli's Principle

Hund's Rule

Periodic trends

Atomic/Ionic Radius

Ionization Energy

Electron Affinity

Electronegativity

Polarizability

Shielding Effect in electrons

CONCEPT MAP

CHEMICAL FORMULA: WRITING AND NAMING

CHEMICAL FORMULA

•  representaLon used to denote a molecule or a formula unit of a pure substance.

•  It indicates the relaLve amounts of atoms of each element in a compound.

•  It consists of the symbols of the elements composing the pure substance and subscripts denoLng the relaLve number of atoms of each element in a molecule or formula unit of the compound.

OXIDATION NUMBER

•  apparent charge of an atom in a compound assuming that electrons are transferred from one atom to another.

•  This set of whole numbers (which may be posiLve or negaLve) are used in predicLng the formulas of compounds, classifying them as ionic or covalent, comparing their chemical properLes, and describing chemical reacLons.

Usual OxidaLon Number (Main Group Elements)

GROUP NO. OXIDATION NUMBERS EXAMPLES

1 +1 Na+1, Li+1

2 +2 Mg2+, Ba2+

13 +3 Al3+, B3+

14 +4/-‐4 C4+, C4-‐

15 -‐3 N3-‐, P3-‐

16 -‐2 O2-‐, S2-‐

17 -‐1 F-‐1, Cl-‐1

OXIDATION NUMBER FORMULA NAME OF THE ION Cr2+ chromous / chromium (ii) Cr3+ chromic / chromium (iii) Mn2+ manganese (ii) Mn4+ manganese (iv) Fe2+ ferrous / iron (ii) Fe3+ ferric / iron (iii) Co2+ cobaltous / cobalt (ii) Co3+ cobalLc / cobalt (iii) Hg22+ mercurous / mercury (i) Hg2+ mercuric / mercury (ii) Sn2+ stannous / Ln (ii) Sn4+ stannic / Ln (iv) Pb2+ plumbous / lead (ii) Pb4+ plumbic / lead (iv)

Common Polyatomic Ions Polyatomic Ions Name

NH4+ ammonium

C2H3O2-‐ acetate

ClO-‐ hypochlorite

ClO2-‐ chlorite

ClO3-‐ chlorate

ClO4-‐ perchlorate

CN-‐ cyanide

OH-‐ hydroxide

HCO3-‐ bicarbonate

IO3-‐ iodate

NO2-‐ nitrite

NO3-‐ nitrate

MnO4-‐ permanganate

HSO4-‐ bisulfate

HSO3-‐ bisulfite

Polyatomic Ions Name

CNO-‐ cyanate

CNS-‐ thiocyanate

CO32-‐ carbonate

CrO42-‐ chromate

Cr2O72-‐ dichromate

C2O42-‐ oxalate

SO32-‐ sulfite

SO42-‐ sulfate

S2O32-‐ thiosulfate

HPO42-‐ biphosphate

SiO32-‐ silicate

ZnO22-‐ zincate

PO33-‐ phosphite

PO43-‐ phosphate

P2O74-‐ pyrophosphate

Rule in assigning oxidaLon number

•  The oxidaLon number of an element in the free or uncombined state is always zero. Example: Cu0, Si0, Mg0

•  The oxidaLon number of a monatomic ion is the same as the charge of the ion.

•  The algebraic sum of the oxidaLon numbers for all the atoms in the formula of a compound is zero. Example: MgBr2 (+2) + (2)(-‐1) = 0

•  The sum of the oxidaLon numbers of atoms in a polyatomic ion must equal the charge of the ion.

Rules in WriLng Chemical Formula •  Metals, nonmetals, and inert gases have their formulas

idenLcal to their symbols. Examples: Calcium, Ca Magnesium, Mg

•  AcLve gaseous elements are wri^en as diatomic molecules (they contain two atoms). Examples: Oxygen, O2 Hydrogen, H2

•  In wriLng formulas for compounds, the symbol of the posiLve element is wri^en first, followed by the symbol of the negaLve element. The sum of the oxidaLon numbers must be equal to zero so that the compound is neutral. Examples: KBr (+1) + (-‐1) = 0 CaCl2 (+2) + 2 (-‐1) = 0

Naming Of Compounds

Binary compounds •  Binary salts – made up of one metallic element and one nonmetallic element. –  For compounds that contain metals with fixed oxidaLon numbers, the name of the caLon is given first followed by the name of the anion, which ends in –IDE.

Examples: NaI sodium iodide MgBr2 magnesium bromide K3N potassium nitride

–  For compounds containing metals with variable oxidaLon numbers, there are two ways of naming them:

•  Old or Classical Method The name of the metal ends with –ous or –ic while the name of the nonmetal ends with –ide. Examples: Cu3B cuprous boride SnS2 stannic sulfide

•  Stock Method The name of the metal is followed by a Roman numeral indica&ng its oxida&on number. This is followed by the name of the nonmetal which ends in –ide. Examples: PbBr2 Lead (II) bromide AuCl3 Gold (III) chloride

Binary Acids •  contain hydrogen and one nonmetallic element.

•  Dry form – The word hydrogen is followed by the name of the nonmetal which ends in –ide. Examples: HCl hydrogen chloride

HBr hydrogen bromide

•  Aqueous form – The prefix hydro-‐ is followed by the name of the nonmetal which ends in –ic acid. Examples: HCl(aq) hydrochloric acid HBr(aq) hydrobromic acid

•  Binary compounds with two nonmetals are named by naming the two elements in the order they appear in the formula. Greek prefixes are used to indicate the number of each element in the compound. The name of the second element ends in –ide. The prefix mono-‐ is not used for the first element. Prefixes: 1 – mono 4 – tetra 7 – hepta 10-‐deca 2 – di 5 – penta 8 – octa 3 – tri 6 – hexa 9 – nona Examples: N2O dinitrogen monoxide CO2 carbon dioxide ICl3 iodine trichloride

•  Binary hydrides contain a metal and hydrogen. The name of the metal comes first followed by the word hydride. Examples: RbH rubidium hydride CsH cesium hydride

TERNARY COMPOUNDS

Ternary acids are named in two ways: •  Dry form – The word hydrogen is followed by the name of the anion. Examples: H2CO3 hydrogen carbonate H2SO3 hydrogen sulfite H2PO4 hydrogen phosphate

•  Aqueous form – Take the root form of the anion’s name. Change –ate endings with –ic, and change –ite endings with –ous. Add the word acid. Examples: H2CO3 carbonic acid H2SO3 sulfurous acid H3PO4 phosphoric acid

•  Ternary salts – Name the caLon and the anion in order. Examples: NaNO3 sodium nitrate FeCO3 iron (II) carbonate

CHEMICAL BONDING

Types of Bonds

Ionic Bond Covalent Bond

r e su l t s f r om e l e c t r o s t aL c a^racLons among ions, which are formed by the transfer of one or more electrons from one atom to another.

r e s u l t s f r o m sharing one or more e le c t ron pairs between two atoms.

90

Lewis Dot Formulas of Atoms

Lewis dot formulas or Lewis dot representa&ons

•  are a convenient bookkeeping method for tracking valence electrons.

Valence electrons •  are those electrons that are transferred

or involved in chemical bonding. •  They are chemically important.

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LEWIS DOT SYMBOLS of Main Group Elements. Elements within a group have the same number of valence electrons and idenLcal Lewis dot symbols.

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Theories of Covalent Bonding

1. Valence Bond (VB) Theory and Orbital HybridizaLon

2. Valence Shell Electron Pair Repulsion Theory (VSEPR)

3. Molecular Orbital Theory (MO)

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VALENCE BOND THEORY Basic Principle •  A covalent bond forms when the orbitals of two atoms overlap and the overlap region, which is between the nuclei, is occupied by a pair of electrons.

94

•  A set of overlapping orbitals has a maximum of two electrons that must have opposite spins.

•  The greater the orbital overlap, the stronger (more stable) the bond.

•  There is a hybridizaLon of atomic orbitals to form molecular orbitals.

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•  OCTET RULE – an atom will form covalent bonds to achieve a complement of eight valence electrons.

–  -‐ majority of main group elements followed octet rule in forming bonds with other elements

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•  EXPANDED OCTET – an atom will form covalent bonds to achieve a complement of more than eight valence electrons ( 10 or 12 valence electrons).

Example :

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•  INEXPANDED OCTET – an atom will form covalent bonds to achieve a complement of less than eight valence electrons ( 2 or 6 valence electrons).

Example :

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Drawing / Wri&ng Lewis Structures STEP 1: Count the t o t a l v a l e n c e e l e c t rons i n t he molecule or ion.

–  Sum the number of valence electrons for each element in a molecule.

–  For ions, add or subtract v a l en ce e l e c t r on s t o account for the charge.

STEP 2: Draw the “skeletal structure” of the molecule.

–  The element wri^en first in the formula is usually the central atom, unless it is hydrogen.

– Usually, the central atom is the LEAST ELECTRONEGATIVE.

STEP 3: Place single bonds between all connected atoms in the s t ruc ture by d r a w i n g l i n e s between them.

–  A single line represents a bonding pair.

STEP 4: Place the remaining valence electrons not accounted for on individual atoms un(l the octet rule is sa(sfied. Place electrons as lone pairs whenever possible.

–  Place electrons first on outer atoms, then on central atoms.

STEP 5: Create mul(ple bonds by shiSing lone pairs into bonding posi(ons as needed for any atoms that do not have a fu l l octet o f va lence electrons.

– Correctly choosing which atoms to form mulLple bonds between comes from experience.

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Predic&ng the Number of bonds in the Lewis Structure ( S – Rule)

N -‐ A = S rule Where: N = number of electrons needed to achieve a noble

gas configuration. •  N usually has a value of 8 for representative

elements. •  N has a value of 2 for H atoms.

A = number of electrons available in valence shells of the atoms.

•  A is equal to the periodic group number for each element.

•  A is equal to 8 for the noble gases. S = number of electrons shared in bonds. (A – S) = number of electrons in unshared, lone, pairs.

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NOTES: •  For ions we must adjust the number of electrons available, A. – Add one e-‐ to A for each nega&ve charge. – Subtract one e-‐ from A for each posi&ve charge.

•  The central atom in a molecule or polyatomic ion is determined by: – The atom that requires the largest number of electrons to complete its octet goes in the center.

– For two atoms in the same periodic group, the less electronegaLve element goes in the center.

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Formal Charges

•  Formal charges are analogous to oxidaLons numbers: – They are not actual charges

•  They keep track of electron ownership

Formal Charge = No. of Valence e- – no. of bonds – no. of non-‐bonding e-‐

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Example 1 : Draw the Lewis structure of OF2.

Step 1: Count the total valence electrons in the molecule or ion.

Step 2: Draw the “skeletal structure” of the molecule.

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STEP 3: Place single bonds between all connected atoms in the structure by drawing lines between them.

STEP 4: Place the remaining valence electrons not accounted for on individual atoms unLl the octet rule is saLsfied. Place electrons as lone pairs whenever possible.

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STEP 5: Create mulLple bonds by shi�ing lone pairs into bonding posiLons as needed for any atoms that do not have a full octet of valence electrons.

MulLple bonds are not required for OF2, as the octet rule is saLsfied for each atom.

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The conceptual steps from molecular formula to the hybrid orbitals used in bonding.

Molecular formula

Lewis structure

GEOMETRY (Electronic/Ideal/e-‐ group arrangement and Molecular

Shape

Hybrid orbitals

Step 1 Step 2 Step 3

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Valence Shell Electron Pair Repulsion (VSEPR) Theory •  The IDEAL/Electronic GEOMETRY of a molecule is determined by the way the electron pairs orient themselves in space •  The orientaLon of electron pairs arises from electron repulsions •  The electron pairs spread out so as to minimize repulsion

107

112

SUMMARY OF IDEAL GEOMETRY

linear trigonal planar tetrahedral

trigonal bipyramidal

octahedral

•  The MOLECULAR GEOMETRY of a molecule is determined by the way the electron pairs orient themselves with respect to their relaLve strength to repel each other. •  The orientaLon of electron pairs arises from electron repulsions

•  The electron pairs spread out so as to minimize repulsion

NOTE: DECREASING repulsion, DECREASING bond angle

LP-‐LP > LP – BP > BP – BP LP – lone (non-‐bonding) pair ; BP – bonding pair

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Species Type Structure Descrip&on Example Bond Angles

AX3

Trigonal Planar

BF3

120o

AX2E

Bent

NO2-‐1

< 120o

Molecular Geometry 3 ELECTRON PAIRS

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Species Type Structure Descrip&on Example Bond Angles

AX4

Tetrahedral

CH4

109.5o

AX3E

Trigonal pyramidal

NH3

< 109.5o

AX2E2

Bent

H2O

< 109.5o

4 ELECTRON PAIRS

–  Methane has four equal bonds, so the bond angles are equal.

–  HOW does the lone pair of ammonia affect its geometry?

–  The bond angles in oxygen are even smaller.

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The conceptual steps from molecular formula to the hybrid orbitals used in bonding.

Molecular formula

Lewis structure

GEOMETRY (Electronic/Ideal/

e-‐ group arrangement and Molecular Shape

HYBRID ORBITALS

Step 1 Step 2 Step 3

Hybrid Orbitals

•  The number of hybrid orbitals obtained equals the number of atomic orbitals mixed.

•  The type of hybrid orbitals obtained varies with the types of atomic orbitals mixed.

Types of Hybrid Orbitals

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Valence Bond : Hybrid Orbitals

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Polar Molecules: The Influence of Molecular Geometry •  Molecular geometry affects molecular polarity.

–  Due to the effect of the bond dipoles and how they either cancel or reinforce each other.

•  Polar Molecules must meet two requirements: 1.  One polar bond or one lone pair of electrons

on central atom. 2.  Neither bonds nor lone pairs can be

symmetrically arranged that their polariLes cancel.

Stoichiometry

Molar Masses of Some Substances

Note: 1 amu = 1.66 x 10-‐24 g

Important Terms in Stoichiometry

Avogadro’s Number, NA •  Number of atoms of an element in a sample whose mass is numerically

equal to the mass of a single atom

1 mole = 6.022 x 1023 parLcles ParLcles = (atoms, ions or molecules

Mole, n •  Avogadro’s number of items

n = mass / molar mass

Chemical Formulas

•  Empirical Formula •  Simplest formula of a compound

•  Molecular Formula •  True formula of a compound •  Whole-‐number mulLple of the Empirical Formula

Note: Empirical formula can be the molecular formula but Molecular may not be the empirical formula

Steps in Determining the Empirical and Molecular Formula of a compound

% Mass of each element

MASS of each element

MOLE of each element

IdenLfy the SMALLEST MOLE of

the element

Divide the MOLES of each element by the SMALLEST MOLE of

the element

MOLE RATIO of each element

Determine the Empirical

Formula Weight (EFW)

Divide the given MOLAR MASS by the EFW

EMPIRICAL FORMULA (simplest Formula)

MOLECULAR FORMULA

KNOW COMMON FRACTIONS IN DECIMAL FORM*

*Page69 Whi^ens (9th ed)

Decimal equivalent

Frac&on Mul&ply by

0.50 1/2 2

0.33 1/3 3

0.67 2/3 3

0.25 ¼ 4

0.75 ¾ 4

0.20 1/5 5

WriLng Chemical EquaLons

1. Write a skeleton equaLon for the reacLon. 2. Indicate the physical state of each reactant

and product. 3. Balance the equaLon

–  Only the coefficients can be changed; subscripts are fixed by chemical nature of the reactants and products

–  It is best to balance atoms that appear only once on each side of the equaLon first

aA(aq) + bB(l) à cC(aq) + dD(s)↓ + eE(g)↑

where: a, b, c, d, e -‐ coefficients of balance equaLon A, B -‐ reactants C, D, E -‐ products ↓ -‐ precipitate formed ↑ -‐ gas formed subscript -‐ state/phase of the reactant & pdt

Mole Method*

Use Molar Use Molar mass (g/mol) mass (g/mol) of Compound A of compound B Use mole ra&o of B and A Use mole ra&o of A and B

*Taken from General Chemistry the essenLal concepts by Raymond Chang Fig 3.8 p 79

MASS (g) of Compound A

MOLE of Compound A

MASS (g) of Compound B

MOLE of Compound B

Approach to LimiLng Reactant Problems

1. Calculate the amount of product that will form if the first reactant were completely consumed.

2. Repeat the calculaLon for the second reactant in the same way.

3. Choose the smaller amount of product and relate it to the reactant that produced it. This is the limiLng reactant and the resulLng amount of product is the theore(cal yield.

4. From the theoreLcal yield, determine how much of the reactant in excess is used, and subtract from the starLng amount.

review in chm11-3/2 concepts

Documents

charge x mass charge

atomic number x

mass of c atom

unit mass

lowest mass

mass number z

e charge

number of p atomic number