Atoms, Molecules and Periodic Table

Copyright © The McGraw-Hill Companies, Inc.  Permission required for reproduction or display.

Dalton’s Atomic Theory (1808)1. Elements are composed of extremely small particles

called atoms.

2. All atoms of a given element are identical, having the same size, mass and chemical properties. The atoms of one element are different from the atoms of all other elements.

3. Compounds are composed of atoms of more than one element.

4. A chemical reaction involves only the separation, combination, or rearrangement of atoms; it does not result in their creation or destruction.

John Dalton• 1800 -Dalton proposed a modern atomic model

based on experimentation not on pure reason.

• All matter is made of atoms.• Atoms of an element are identical.• Each element has different atoms.• Atoms of different elements combine

in constant ratios to form compounds.• Atoms are rearranged in reactions.

• His ideas account for the law of conservation of mass (atoms are neither created nor destroyed) and the law of constant composition (elements combine in fixed ratios).

Dalton’s Atomic Theory

Law of Multiple Proportions

8 X2Y16 X 8 Y+

Law of Conservation of Mass

Ernest Rutherford (movie: 10 min.)

Most particles passed through. So, atoms are mostly empty.

Some positive -particles deflected or bounced back!

Thus, a “nucleus” is positive & holds most of an atom’s mass.

Radioactive substance path of invisible

-particles

• Rutherford shot alpha () particles at gold foil.

Lead block

Zinc sulfide screen

Thin gold foil

atomic radius ~ 100 pm = 1 x 10-10 m

nuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m

Rutherford’s Model of the Atom

“If the atom is the Houston Astrodome, then the nucleus is a marble on the 50-yard line.”

mass p ≈ mass n ≈ 1840 x mass e-

Atomic number (Z) = number of protons in nucleus

Mass number (A) = number of protons + number of neutrons

= atomic number (Z) + number of neutrons

Isotopes are atoms of the same element (X) with different numbers of neutrons in their nuclei

XAZ

H11 H (D)2

1 H (T)31

U23592 U238

92

Mass Number

Atomic NumberElement Symbol

Atomic number, Mass number and Isotopes

The Isotopes of Hydrogen

6 protons, 8 (14 – 6) neutrons, 6 electrons

6 protons, 5 (11 – 6) neutrons, 6 electrons

How many protons, neutrons, and electrons are in C14

6 ?

How many protons, neutrons, and electrons are in C11

6 ?

Properties of Waves

Wavelength () is the distance between identical points on successive waves.Amplitude is the vertical distance from the midline of a wave to the peak or trough.Frequency () is the number of waves that pass through a particular point in 1 second (Hz = 1 cycle/s).The speed (u) of the wave = x

1. e- can only have specific (quantized) energy values

2. light is emitted as e- moves from one energy level to a lower energy level

Bohr’s Model of the Atom (1913)

En = -RH ( )1n2

n (principal quantum number) = 1,2,3,…

RH (Rydberg constant) = 2.18 x 10-18J

E = h

E = h

The Planck constant was first discovered as the proportionality constant between the energy (E) of a photon and the frequency of its associated electromagnetic wave (ν) length. This relation between the energy and frequency is called the Planck relation or the Planck–Einstein equation:

E=hvh=6.626×10−34

Since the frequency ν, wavelength λ, and speed of light c are related by λν = c, the Planck relation can also be expressed as                            

c=3x108  speed of the

light                                                                  

Bohr’s model

There are 2 types of spectra: continuous spectra & line spectra. It’s when electrons fall back down that they release a photon. These jumps down from “shell” to “shell” account for the line spectra seen in gas discharge tubes (through spectroscopes).

• Electrons orbit the nucleus in “shells”•Electrons can be bumped up to a higher

shell if hit by an electron or a photon of light.

Schrodinger Wave Equation

is a function of four numbers called quantum numbers (n, l, ml, ms)

Quantum Numbers of Electrons in Atoms

Name Symbol Permitted Values Property

principal n positive integers(1,2,3,…) orbital energy (size)

angular momentum

l integers from 0 to n-1 orbital shape (The l values 0, 1, 2, and 3 correspond to s, p, d, and f orbitals, respectively.)

magnetic ml integers from -l to 0 to +l orbital orientation

spin ms +1/2 or -1/2 direction of e- spin

The Pauli Exclusion Principle - No two electrons in the same atom can have the same four q.n. Since the first three q.n. define the orbital, this means only two electrons can be in the same orbital and they must have opposite spins.

Schrodinger Wave Equation is a function of four numbers called quantum numbers (n, l, ml, ms)

principal quantum number n

n = 1, 2, 3, 4, ….

n=1 n=2 n=3

distance of e- from the nucleus

quantum numbers: (n, l, ml, ms)

angular momentum quantum number l

for a given value of n, l = 0, 1, 2, 3, … n-1 from 0 to n-1

n = 1, l = 0n = 2, l = 0 or 1

n = 3, l = 0, 1, or 2

Shape of the “volume” of space that the e- occupies

l = 0 s orbitall = 1 p orbitall = 2 d orbitall = 3 f orbital

Schrodinger Wave Equation

l = 0 (s orbitals)

l = 1 (p orbitals)

l = 2 (d orbitals)Clover leaf

quantum numbers: (n, l, ml, ms)

magnetic quantum number ml

for a given value of lml = -l, …., 0, …. +l

From –l to 0 then to + l

orientation of the orbital in space

if l = 1 (p orbital), ml = -1, 0, or 1if l = 2 (d orbital), ml = -2, -1, 0, 1, or 2

Schrodinger Wave Equation

ml = -1, 0, or 1 3 orientations is space

ml = -2, -1, 0, 1, or 2 5 orientations is space

(n, l, ml, ms)

spin quantum number ms

ms = +½ or -½

Schrodinger Wave Equation

ms = -½ms = +½

Schrodinger Wave Equationquantum numbers: (n, l, ml, ms)

Shell – electrons with the same value of n

Subshell – electrons with the same values of n and l

Orbital – electrons with the same values of n, l, and ml

How many electrons can an orbital hold?

If n, l, and ml are fixed, then ms = ½ or - ½

= (n, l, ml, ½)or= (n, l, ml, -½)

An orbital can hold 2 electrons

How many 2p orbitals are there in an atom?

2p

n=2

l = 1

If l = 1, then ml = -1, 0, or +1

3 orbitals

How many electrons can be placed in the 3d subshell?

3d

n=3

l = 2

If l = 2, then ml = -2, -1, 0, +1, or +2

5 orbitals which can hold a total of 10 e-

Energy of orbitals in a single electron atom

Energy only depends on principal quantum number n

En = -RH ( )1n2

n=1

n=2

n=3

The most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins (Hund’s rule).

s has one orbital, p has 3 orbitals, d has 5 orbitals.

Each orbitals can hold upto two electrons

Order of orbitals (filling) in multi-electron atom

1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s

Electron configuration is how the electrons are distributed among the various atomic orbitals in an atom.

1s1

principal quantumnumber n

angular momentumquantum number l

number of electrons in the orbital or subshell

Orbital diagram

H

1s1

34

Isoelectronics, the number of electrons on each ions the same In an isoelectronic series of ions, the nuclear charge increases with increasing atomic number and draws the electrons inward with greater force. The ion with fewest proton produces the weakest attractive force on the electrons and thus has the largest size

GROUND STATE VS EXCITED STATE of Electron energy state.

The energy associated to an electron is that of its orbital. The energy of a configuration is often approximated as the sum of the energy of each electron, neglecting the electron-electron interactions.

The configuration that corresponds to the lowest electronic energy is called the ground state. Any other configuration is an excited state.

36

What is the electron configuration of Mg?

Mg 12 electrons

1s < 2s < 2p < 3s < 3p < 4s

1s22s22p63s2 2 + 2 + 6 + 2 = 12 electrons

Abbreviated as [Ne]3s2 [Ne] 1s22s22p6

What are the possible quantum numbers for the last (outermost) electron in Cl?

Cl 17 electrons 1s < 2s < 2p < 3s < 3p < 4s

1s22s22p63s23p5 2 + 2 + 6 + 2 + 5 = 17 electrons

Last electron added to 3p orbital

n = 3 l = 1 ml = -1, 0, or +1 ms = ½ or -½

Outermost subshell being filled with electrons

Paramagneticunpaired electrons

2p

Diamagneticall electrons paired

2p

THE MODERN PERIODIC TABLE

The Periodic Table

Copyright © The McGraw-Hill Companies, Inc.  Permission required for reproduction or display.

42

ns1

ns2

ns2

np1

ns2

np2

ns2

np3

ns2

np4

ns2

np5

ns2

np6

d1

d5 d10

4f

5f

Ground State Electron Configurations of the Elements

43

Classification of the Elements

Metals

Nobel Gases

44

Electron Configurations of Cations and Anions

Na [Ne]3s1 Na+ [Ne]

Ca [Ar]4s2 Ca2+ [Ar]

Al [Ne]3s23p1 Al3+ [Ne]

Atoms lose electrons so that cation has a noble-gas outer electron configuration.

H 1s1 H- 1s2 or [He]

F 1s22s22p5 F- 1s22s22p6 or [Ne]

O 1s22s22p4 O2- 1s22s22p6 or [Ne]

N 1s22s22p3 N3- 1s22s22p6 or [Ne]

Atoms gain electrons so that anion has a noble-gas outer electron configuration.

Of Representative Elements

45

+1

+2

+3 -1-2-3

Cations and Anions Of Representative Elements

46

Effective nuclear charge (Zeff) is the “positive charge” felt by an electron.

Na

Mg

Al

Si

11

12

13

14

10

10

10

10

1

2

3

4

186

160

143

132

ZeffCoreZ Radius (pm)

Zeff = Z - 0 < < Z ( = shielding constant)

Zeff Z – number of inner or core electrons

47

Effective Nuclear Charge (Zeff)

increasing Zeff

incr

easi

ng Z

eff

48

Atomic Radii

metallic radius covalent radius

49

50

Cation is always smaller than atom from which it is formed.Anion is always larger than atom from which it is formed.

51

Ionization energy is the minimum energy (kJ/mol) required to remove an electron from a gaseous atom in its ground state.

I1 + X (g) X+(g) +

e-

I2 + X+(g) X2+

(g) + e-

I3 + X2+(g) X3+

(g) + e-

I1 first ionization energy

I2 second ionization energy

I3 third ionization energy

I1 < I2 < I3

52

General Trends in First Ionization Energies

Increasing First Ionization Energy

Incr

ea

sing

Firs

t Io

niz

atio

n E

ner

gy

53

Electron affinity is the negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion.

X (g) + e- X-(g)

F (g) + e- X-(g)

O (g) + e- O-(g)

H = -328 kJ/mol EA = +328 kJ/mol

H = -141 kJ/mol EA = +141 kJ/mol

54

55

Types of Elements

Metals, Nonmetals, metalloids, transition elements,

Metals, found on both the left side and the in the middle of the table, include active metals, transition metals and the lanthanide and actinide series of elements.

They are shiny solids except for mercury. They have ability to deform without breaking which is called malleability. They have high melting points, and density.

They have low electronegativity, large atomic radius, low ionization energy. They easily give up electrones.

57

Nonmetals are found on the upper right side of the table. Nonmetals are generally brittle in the solid state and show little or no metallic luster. They have high ionization energy, electron affinity and electro negativity. They have small atomic radii.

Metalloids are also called semi metals. They posses characteristics of metals and nonmetals. For example Silicon has a metallic luster but is brittle and a poor electric conductor

58

Group 1A Elements (ns1, n 2)

M M+1 + 1e-

2M(s) + 2H2O(l) 2MOH(aq) + H2(g)

4M(s) + O2(g) 2M2O(s)

Incr

easi

ng r

eact

ivity

59

Group 2A Elements (ns2, n 2)

60

Group 2A Elements (ns2, n 2)

M M+2 + 2e-

Be(s) + 2H2O(l) No ReactionIn

crea

sing

rea

ctiv

ity

Mg(s) + 2H2O(g) Mg(OH)2(aq) + H2(g)

M(s) + 2H2O(l) M(OH)2(aq) + H2(g) M = Ca, Sr, or Ba

61

Group 3A Elements (ns2np1, n 2)

4Al(s) + 3O2(g) 2Al2O3(s)

2Al(s) + 6H+(aq) 2Al3+

(aq) + 3H2(g)

62

Group 3A Elements (ns2np1, n 2)

63

Group 4A Elements (ns2np2, n 2)

Sn(s) + 2H+(aq) Sn2+

(aq) + H2 (g)

Pb(s) + 2H+(aq) Pb2+

(aq) + H2 (g)

64

Group 4A Elements (ns2np2, n 2)

65

Group 5A Elements (ns2np3, n 2)

N2O5(s) + H2O(l) 2HNO3(aq)

P4O10(s) + 6H2O(l) 4H3PO4(aq)

66

Group 5A Elements (ns2np3, n 2)

67

Group 6A Elements (ns2np4, n 2) (oxygen family=chalcogen)

SO3(g) + H2O(l) H2SO4(aq)

68

Group 6A Elements (ns2np4, n 2)

69

Group 7A Elements (ns2np5, n 2)

X + 1e- X-1

X2(g) + H2(g) 2HX(g)

Incr

easi

ng r

eact

ivity

70

Group 7A Elements (ns2np5, n 2)

71

Group 8A Elements (ns2np6, n 2)

Completely filled ns and np subshells.Highest ionization energy of all elements.No tendency to accept extra electrons.

72

Compounds of the Noble Gases

A number of xenon compounds XeF4, XeO3, XeO4, XeOF4 exist.A few krypton compounds (KrF2, for example) have been prepared.

73

Properties of Oxides Across a Period

basic acidic

Period

Group

Alkali M

etal

Noble G

as

Halogen

Alkali E

arth Metal

The Modern Periodic Table

A molecule is an aggregate of two or more atoms in a definite arrangement held together by chemical forces

H2 H2O NH3 CH4

A diatomic molecule contains only two atoms

H2, N2, O2, Br2, HCl, CO

A polyatomic molecule contains more than two atoms

O3, H2O, NH3, CH4

diatomic elements

An ion is an atom, or group of atoms, that has a net positive or negative charge.

cation – ion with a positive chargeIf a neutral atom loses one or more electronsit becomes a cation.

anion – ion with a negative chargeIf a neutral atom gains one or more electronsit becomes an anion.

Na 11 protons11 electrons Na+ 11 protons

10 electrons

Cl 17 protons17 electrons Cl–

17 protons18 electrons

A monatomic ion contains only one atom

A polyatomic ion contains more than one atom

Na+, Cl-, Ca2+, O2-, Al3+, N3-

OH-, CN-, NH4+, NO3

-

Common Ions Shown on the Periodic Table

13 protons, 10 (13 – 3) electrons

34 protons, 36 (34 + 2) electrons

How many protons and electrons are in ?Al2713

3+

How many protons and electrons are in ?Se7834

2-

A molecular formula shows the exact number of atoms of each element in the smallest unit of a substance

An empirical formula shows the simplest whole-number ratio of the atoms in a substance

H2OH2O

molecular empirical

C6H12O6 CH2O

O3 O

N2H4 NH2

Ionic compounds consist of a combination of cations and an anions

• The formula is usually the same as the empirical formula

• The sum of the charges on the cation(s) and anion(s) in each formula unit must equal zero

The ionic compound NaCl

Formula of Ionic Compounds

Al2O3

2 x +3 = +6 3 x -2 = -6

Al3+ O2-

CaBr2

1 x +2 = +2 2 x -1 = -2

Ca2+ Br-

Na2CO3

1 x +2 = +2 1 x -2 = -2

Na+ CO32-

The most reactive metals (green) and the most reactive nonmetals (blue) combine to form ionic compounds.

Chemical Nomenclature

• Ionic Compounds– Often a metal + nonmetal– Anion (nonmetal), add “ide” to element name

BaCl2 barium chloride

K2O potassium oxide

Mg(OH)2 magnesium hydroxide

KNO3 potassium nitrate

• Transition metal ionic compounds– indicate charge on metal with Roman numerals

FeCl2 2 Cl- -2 so Fe is +2 iron(II) chloride

FeCl3 3 Cl- -3 so Fe is +3 iron(III) chloride

Cr2S3 3 S-2 -6 so Cr is +3 (6/2) chromium(III) sulfide

• Molecular compounds− Element furthest to the left in a

period and closest to the bottom of a group on periodic table is placed first in formula

− If more than one compound can be formed from the same elements, use prefixes to indicate number of each kind of atom

− Last element name ends in ide

HI hydrogen iodide

NF3 nitrogen trifluoride

SO2 sulfur dioxide

N2Cl4 dinitrogen tetrachloride

NO2 nitrogen dioxide

N2O dinitrogen monoxide

Molecular Compounds

1. Which of the following increases with increasing atomic number within a family (group) on the periodic table?

a. electro negativityb. electron affinityc. atomic radiusd. ionization energy

2. Silicon has a silvery luster at room temperature. Silicon is brittle, and does not conduct heat or electricity well. Based on its position in the periodic table, silicon is most likely a:

a. Nonmetalb. Metalloidc. Metald. Chalcogen (Oxygen group)

3. A natural sample of carbon contains 99% of 12C. How many moles of 12C are likely to be fund in a 48.5 gram samle of carbon obtained from nature?

a. 1b. 2c. 4d. 12e. 49.5

91

4. Which of the following is the correct electron configuration for Zn2+

a. 1s22s22p63s23p64s03d10

b. 1s22s22p63s23p64s23d8

c. 1s22s22p63s23p64s23d10

d. 1s22s22p63s23p64s03d8

92

5. Which of the following quantum number sets describes a possible element?

a. n=2; l=2; m1=1; ms=+1/2b. n=2; l=1; m1=-1; ms=+1/2c. n=2; l=0; m1=-1; ms=-1/2d. n=2; l=0; m1=-1; ms=-1/2

93

6. What is the maximum number of electrons allowed in a single atomic energy level in terms of the principal quantum number of n?

a. 2nb. 2n+2c. 2n2

d. 2n2+2

94

8. An electron returns from an excited state to its ground state, emitting a photon at 500nm. What would be the magnitude of the energy change if this process were repeated such that a mole of these photons were emitted?

a. 3.98x10-19 Jb. 3.98x10-21 Jc. 2.39x105 Jd. 2.39x10-3 J

95

Answer 8

E=(6.26x10-34)x((3X108)/500 x10-9m)E=(6.26x10-34)x(6x1014)E=3.98x10-19J this is per photonE=3.98x10-19x 6.022x1023=2.39x105J

96

Q.9 Which of the following statement is NOT true of an electrons grounds state?

a. The electro is at its lowest possible energy level.

b.The electron is in a quantized energy level,

c. The electon is traveling along its smallest possible orbital rdius

d. The electron is static

97

Q. 10 Which of the following experimental conditions would NOT excite an electron out of the ground state?

a. Radiationb. High temperaturec. High pressured. None of the above

Q11. What determines the length of an elements atomic radius?

I. the number of valence electrons

II. The number of electron shells

III. The number of neutrons in the nucleus

a. I onlyb. II onlyc. I and II onlyd. I, II and III

99

Q. 12. How many valence electrons are present in elements third period?

a. 2b. 3c. The number decreases as the

atomic number increasesd. The number increases as the

atomic number increases