Chapter 2

Chapter 2

Atoms

What is Chemistry?

• The study of matter and its properties and transformations

What is Matter?• Anything that has mass and volume

– Mass = the amount of a substance, measured in grams, g

– Volume = the space occupied by a substance, measured in cm3, mL (milliliters) or L (Liters)

Brief History of Chemistry

• Ideas about matter date back to ancient Greece 2000 years ago

• 2 schools of thoughtDemocritus – all matter made of tiny, indivisible

particles called atoms (from Greek word atomos = uncuttable)

Aristotle – matter is continuous, it is infinitely divisible

• Aristotle’s ideas dominated for almost 2000 yrs

So, who was right?

• Today we know that Democritus was right• Atoms are the basic building blocks of matter• we will discuss evidence for the existence of

atoms later

Why do we believe in atoms?

• First atomic theory based on scientific evidence proposed in 1808 by English chemist John Dalton (1766 -1844)

• Theory based on three scientific laws discovered in late 1700s, early1800s

• Law of Conservation of Mass (Antoine Lavoisier (1743-1794))– Matter cannot be created or destroyed in an

ordinary chemical reaction

• Law of Constant Composition (Joseph Proust (1754-1826))– no matter where you find a specific compound, it

is always made up of the same proportion of elements by mass elements combine to form compounds in fixed proportions

• Law of Multiple Proportions (John Dalton)– elements always combine to make compounds in

whole number ratios or multiples of whole number ratios, never in fractions

(Not mentioned in book)

Postulates of Dalton’s Atomic Theory

1. All matter is made of tiny, indivisible particles called atoms (in honor of Democritus’ idea)

2. All atoms of a given element have the same properties; atoms of different atoms have different properties

3. Compounds are formed by the chemical combinations of two or more different types of atoms

4. During chemical reactions, atomic arrays are just rearranged into new combinations

Dalton’s Model of the Atom• Atoms are solid, indivisible spheres, like

billiard balls• His model was referred to as the “Billiard Ball”

model• Dalton’s model of the atom was to endure for

almost 100 years, until the discovery of radioactivity and the first subatomic particle (the electron) in the late 1890s.

What are atoms made of?

• First subatomic particle, the electron, discovered by English Physicist J.J. Thomson in 1897.

• The proton was discovered in 1919 by Ernest Rutherford

• The last subatomic particle to be identified was the Neutron in 1932 by James Chadwick.

Properties of Subatomic Particles

ParticleParticle LocationLocationRelative Relative electricaelectrical chargel charge

Actual Actual

Mass (g)Mass (g)

ElectronElectron In space In space surround-surround-ing ing nucleusnucleus

1-1- 9.11 x 109.11 x 10-28-28

ProtonProton In nucleusIn nucleus 1+1+ 1.673 x 101.673 x 10--

2424

NeutronNeutron In nucleusIn nucleus 00 1.675 x 101.675 x 10--

2424

RelativeRelative

Mass Mass (amu)(amu)

__1____1__

18401840

11

11

Terminology for Atomic Structure

Atomic number (Z) – the number of protons in the nucleus of an atom, also the number of electrons as atoms are electrically neutral

Mass number (A) – the number of protons and neutrons in the nucleus of an atom

• Number of neutrons in the nucleus : #neutrons = mass no. – atomic no.

C

Subatomic Particles

6

Determining the number of protons and electrons in an atom from the periodic table

6

Atomic number

Symbol

Atomic number = # protons

= 6

= # electrons

= 6

Carbonn

12.01

C

Subatomic Particles

Determining the number of neutrons in an atom: Mass # - Atomic #

- Must be given the mass number!- Mass number is not the same as the atomic

mass- e.g. Sodium with a mass number of 23 Na

atomic # = 11, 11 protons, 11 e-neutrons = 23 – 11 = 12 neutrons

Discovery Of Isotopes

• After neutrons discovered, it was found that not all atoms of the same element were the same (as Dalton had said)

• Almost every element has examples of atoms that have the same number of protons, but different numbers of neutrons

Isotopes = Atoms that have the same number of protons (atomic number), but different numbers of neutrons (different mass numbers)

Nuclear NotationNuclear Notation Contains the symbol of the Contains the symbol of the

element, the mass number and the element, the mass number and the atomic number.atomic number.

X Massnumber

Atomicnumber

Examples

C

Number of protons = 6Number of electrons = 6Number of neutrons = 12 – 6 = 6

Mass number can also be used at the end of the element’s name e.g. carbon-12

12

6

Atomic number

Mass number

Isotope ExamplesIsotopeIsotope Mass Mass

number number (A)(A)

AtomiAtomic c numbnumber (Z)er (Z)

# # protonprotonss

# # neutronsneutrons

# # electronelectronss

Carbon-Carbon-1313

Cobalt - Cobalt - 5858

Sodium-Sodium-2323

Isotope ExamplesIsotopeIsotope Mass Mass

number number (A)(A)

AtomiAtomic c numbnumber (Z)er (Z)

# # protonprotonss

# # neutronsneutrons

# # electronelectronss

Carbon-Carbon-1313

1313 66 66 77 66

Cobalt - Cobalt - 5858

5858 2727 2727 3131 2727

Sodium-Sodium-2323

2323 1111 1111 1212 1111

Introduction to the Periodic Table

Dmitri Mendeleev (1834-1907)

• When Mendeleev arranged the elements by increasing atomic weight, he noticed a periodic repetition in atomic properties (e.g. density, melting point)

•Because the properties of the atoms were repeated periodically, the table he created was called periodic table

Mendeleev’s 1872 Table

Modern Periodic Table

• After Moseley discovered atomic number, elements were rearranged from increasing atomic weight to increasing atomic number

• Vertical Columns called groups or families. • Horizontal rows called periods.


IA

IIA IIIA IVA VA VIAVIIA

VIIA

Group number

1

2

3

5

4

6

7

Period number

3 Main Categories of Elements

1. Metals – Shiny, good conductors of electricity and heat, tend to have 3or less valence e-, malleable, ductile, located to the left of the stair step on the periodic table

2. Nonmetals – dull, brittle, poor (some non) conductors of electricity and heat, have 4 or more valence e-, located to the right of the stair step on the periodic table

Stair Step Nonmetals

Metals

Metalloids

3 Main Categories of Elements

3. Metalloids – located along the stair step on the periodic table, have properties of both metals and nonmetalse.g. Silicon – is shiny, brittle, semiconductor of electricity


Alkali metals


Alkaline earth metals

Introduction to the Periodic TableHalogens

Introduction to the Periodic TableNoble gases


Transition metals


Lanthanide series

Actinideseries


Metalloids


Main Group

Elements

Periodicity

• Trend within a group of elements or across a period of elements in the periodic table

How are the electrons in an atom arranged?

• Atom is mostly empty space with central dense core called nucleus

• Electrons are located at a distance away and have to be constantly moving to avoid being pulled into the positively charged nucleus

• Because e- are moving, they possess kinetic energy

• In 1913, Niels Bohr discovered

Niels Bohr 1913

• Discovered that only certain values are possible for the energy of the hydrogen electron

• The energy of the electron is quantized only certain values are allowed

Quantized Energy Levels

• The energy levels of all atoms are quantized.

• Electrons are confined to specific regions of space, called principal energy levels or shells

• These energy levels or shells radiate away from the nucleus and given whole integer numbers of 1, 2, 3, 4, etc

• Each energy level can accommodate only a certain number of electrons, given by the formula 2n2

Energy level maximum number of (Shell) n electrons (2n2) 1 2 2 8 3 18 4 32

Energy levels are further divided into sublevels or subshells

• Sublevels are designated by the letters s, p, d and f

n = 1 1 sublevel = 1s n = 2 2 sublevels = 2s and 2p n= 3 3 sublevels = 3s, 3p and 3d n = 4 4 sublevels = 4s, 4p, 4d, and 4f

Within these sublevels, electrons are grouped in orbitals

Orbital = most probable region in space of finding an electron

• According to quantum theory, there is a limit to what we can know about the electron

• Therefore, we can only discuss its location in terms of probability.

• Orbitals are probability maps that have definite shapes and orientations in space

• Each orbital can hold a maximum of 2 electrons

Sublevel designation s, p, d and f also designates the shape of the electron

orbital

S orbitals = spherical an shape

1s 2s 3s

p orbitals

• p orbitals are dumbbell shaped• There are three p orbital shapes

• The s and p types of sublevel

d Orbitals

7 -

f-orbtals

http://www.d.umn.edu/~pkiprof/ChemWebV2/AOs/ao4.html

http://www.d.umn.edu/~pkiprof/ChemWebV2/AOs/ao4.html

Electron Configuration

Electron configuration:Electron configuration: The arrangement of electrons in the extranuclear space (i.e. the empty space surrounding the nucleus).

• The energy of the electrons in an atom is quantizedquantized, which means that an electron in an atom can have only certain allowed energies.

• These allowed energies correspond to specific regions in space surrounding the nucleus called energy levels or shells. Ground-state electron configuration:Ground-state electron configuration: The electron configuration of the lowest energy state of an atom.

Electron configurations

• Tell us the orbital location of an atom’s electrons

• Like an address


Table 2.5 Distribution of Electrons in Shells

Shell

4321

321882

Relativeenergies

of electrons in each shell

lower

higher

Number ofelectrons shell

can hold

Assigning electrons to orbitals

• Orbital filling diagrams use boxes or circles to represent orbitals

• See handout• Rules for filling orbitals

– Bottom up rule – atoms place their electrons in the lowest possible energy orbitals first

– Each orbital can hold a maximum of 2 electrons, which must be spinning in opposite directions

– For p, d and f orbitals, one electron in each orbital before pairing up


Table 2.6 Distribution of Orbitals within Shells

Shell

321

One 3s, three 3p, and five 3d orbitalsOne 2s and three 2p orbitalsOne 1s orbital

Orbitals Contained in Each Shell

4 One 4s, three 4p, five 4d, and seven 4f orbitals

Maximum Numberof Electrons Shell

Can hold

22 + 6 = 82 + 6 + 10 = 182 + 6 + 10 + 14 = 32

Electron ConfigurationFigure 2.13 Energy levels for orbitals through the third shell.


Electron configurations are governed by three rules:Rule 1:Rule 1: Orbitals fill in the order of increasing energy from lowest to highest.– Elements in the first, second, and third periods

fill in the order 1s, 2s, 2p, 3s, and 3p.


Rule 2:Rule 2: Each orbital can hold up to two electrons with spins paired in opposite directions.– With four electrons, the 1s and 2s orbitals are

filled and are written 1s2 2s2.– With an additional six electrons, the three 2p

orbitals are filled and are written either 2px2 2py

2 2pz

2, or they may be written 2p6.


Orbitals have definite shapes and orientations in space

Electron ConfigurationFigure 2.14 The pairing of electron spins.


Rule 3:Rule 3: When there is a set of orbitals of equal energy, each orbital becomes half filled before any of them becomes completely filled.– Example:Example: After the 1s and 2s orbitals are filled, a

5th electron is put into the 2px, a 6th into the 2py, and a 7th into the 2pz. Only after each 2p orbital has one electron is a second added to any 2p orbital.


Orbital box diagramsOrbital box diagrams– A box represents an orbital.– An arrow represents an electron.– A pair of arrows with heads in opposite

directions represents a pair of electrons with paired spins.

ExampleExample: carbon (atomic number 6)

1s2 2s2 2px1 2py

1Expanded:

1s 2s

Electron configuration

Condensed:1s2 2s2 2p22px 2py 2pz


Noble gas notation– The symbol of the noble gas immediately

preceding the particular atom indicates the electron configuration of all filled shells

Example:Example: carbon (atomic number 6)Electron

Configuration(condensed)Orbital box diagram

Noble GasNotation

1s2 2s2 2p2 [He]2s2 2p2

Electron ConfigurationValence shell:Valence shell: The outermost incomplete shell.

Valence electron:Valence electron: An electron in the valence shell.Lewis dot structure: – The symbol of the element represents the nucleus and

filled shells.

– Dots represent valence electrons.

N OB

H

Li Be

Na Mg

He1A 2A 5A 6A 7A 8A3A 4A

Cl

F

S

Ne

ArSiAl P

C


8A

1s7A6A5A4A3A

2B1B8B8B

1A

8B7B

2A

6B5B4B3B

1

2

3

4

5

6

7

1

2

3

4

5

6

7

6

7

Main group elements;s block (2 elements)

Transition elements;d block (10 elements)

Main group elements;p block (6 elements)

1s

Helium is also an s blockelement

3d

4d

5d

3p

4p

6d

2s

3s

4s

5s

6s

7s

2p

5p

6p

7p

4f

5f

Inner transition elements; f block

(14 elements)


Table 2.9 Noble Gas Notation and Lewis dot structures for the Alkali Metals (Group 1A Elements)

Cs

RbCs•

K

Rb•

1A

K•

Na

Na•

Cs

Li•

Rb

Li

K

Na

Li

Element

NobleGas

Notation

Lewisdot

Structure

[He]2s1

[Ne]3s1

[Ar]4s1

[Kr]5s1

[Xe]6s1

3

6.94111

22.99019

39.09837

85.46855

132.91

Periodic Property

• As we have seen, the Periodic Table was constructed on the basis of trends (periodicity) in chemical properties.

• With an understanding of electron configuration, chemists realized that the periodicity of chemical properties could be understood in terms of periodicity in electron configuration.

• The Periodic Table worked because elements in the same column (group) have the same configuration in their outer shells.

• We look at two periodic properties: Atomic size and ionization energy.

Atomic Size

The size of an atom is determined by the size of its outermost occupied orbital.

• Example: The size of a chlorine atom is determined by the size of its three 3p orbitals, the size of a carbon atom is determined by the size of if its three 2p orbitals.

198 pm154 pm

The radius of a chlorineatom is 99 pm

The radius of a carbonatom is 77 pm

Cl Cl C C

Atomic SizeFigure 2.16 Atomic radii of the main-group elements (in picometers).

Ionization Energy

Ionization energy:Ionization energy: The energy required to remove the most loosely held electron from an atom in the gaseous state.– Example: When lithium loses one electron, it

becomes a lithium ion; it still has three protons in its nucleus, but now only two electrons outside the nucleus, and therefore has a positive charge.

Li (g) + Li+(g) + e-

Lithium Lithiumion

ElectronIonizationenergy

energy

Ionization Energy

Ionization energy is a periodic property:

– In general, it increases across a row; valence electrons are in the same shell and subject to increasing attraction as the number of protons in the nucleus increases.

– It increases going up a column; the valence electrons are in lower principle energy levels, which are closer to the nucleus and feel the nuclear charge more strongly.

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Chapter 2