atomic structure the half of knowledge is knowing where to find knowledge. unknown ch2. 1. j.c. rowe...

ATOMIC STRUCTURETHE HALF OF KNOWLEDGE IS KNOWING WHERE TO FIND KNOWLEDGE. UNKNOWN

Ch2. 1. J.C. Rowe

Windsor University School of Medicine

CONCEPT MAPElements,

compounds & mixtures

Separation methods Solutions

Atoms have structure

Atomic nucleus

Electronic structure

Size of atoms

Atoms & ions

Symbols for atoms & ions

Elements, Mixtures and Compounds

Elements: Are materials that could not be separated into simpler things

either physically or chemically Elements contain only one type of atom Iron is an example of an element

Mixture: Composition of two or more substances that can be separated

physically Example of a chemical mixture is the combination of Sand and

Salt

Compound: Combination of two or more elements The units of a compound have different properties

Compounds have fixed compositionsExample – combination of Iron and Sulphur.

Mixtures

• What is a mixture?

• How can it be separated?

• A mixture is a physical blend of two or more substances.

• Their composition varies (Air).

• There are two types of mixtures: homogenous and heterogenous

Homogenous or Heterogenous?

1. Air

2. Salt water

3. Tea

4. Brass

5. Vinegar

6. Hydrogen peroxide

7. Steel

1. Salad dressing

2. Apple

3. Sand

4. Paint

5. Granite

6. Laundry detergent

7. Cereal

Heterogenous Mixtures Is the type of mixture that is not uniform in

composition.

If one portion of the above mixture were sampled, it’s composition would vary.

Homogenous Mixtures Is a type of mixture that has a completely

uniform composition throughout. It’s components are evenly distributed

throughout the sample.

Solutions

• Is the special name that scientists give to homogenous mixtures.

• Solutions may be of gases, liquids or solids.

• An example:

solution of sugar solid in liquid water.

What is a Solution ?Ques:

• A Mixture of 2 or more soluble substances dissolved in a solvent

• Can you name any solutions ?

What is a solution ?

Solutions are a special kind of mixture.

Solution = solvent (major) + solute (minor).

A solvent may be a solid, a liquid or a gas.

When the solvent is a liquid the solution is said to be aqueous (water) or non-aqueous (other than water).

Solution can be homogeneous or

heterogeneous.

Solutions

Some common types of solutionsSome common types of solutions

System Examples System Examples

Gas-gasGas-gas COCO2 2 and O in N (air) and O in N (air)

Liquid-gasLiquid-gas Water vapor in airWater vapor in air

Gas-liquidGas-liquid COCO22 in in HH22O (Soda water)O (Soda water)

Liquid-liquidLiquid-liquid Acetic acid in HAcetic acid in H22O (vinegar)O (vinegar)

Solid-liquidSolid-liquid NaCl in NaCl in HH22O (brine)O (brine)

Solid-solidSolid-solid Cu in Ag (Sterling silver)Cu in Ag (Sterling silver)

Separating Mixtures

If you were given a mixture of iron nails, salt and water…

How would you separate this mixture completely?

Based on which physical properties would you base your method on?

Separating Mixtures

How would you separate the components in tap water?

Distillation

A liquid is boiled to produce vapor that is then condensed again to a liquid

Separating Mixtures

Conditions for solubility

Temperature Solids usually are more soluble in liquids

as the temperature increases; Gases are less soluble in warm water than

cold water; Gases dissolve more in liquids at higher

pressure. Ionic materials dissolve in water Covalent materials dissolve in

covalent liquids

Separating mixtures

Filtration

Evaporation

Distillation/fractional distillation

Gravity separation

chromatography

The nuclear atom

Rutherford showed in 1911 that atoms:

Have a very tiny nucleus at the center Have all their positive charge in the

nucleus Have almost all their mass in the nucleus Have only negative electrical charges

making up the rest of their volume

Model of an Atom

The Rutherford Experiment

Properties of the 3 sub-atomic particles

Name of particle

symbol Relative mass

charge location

Electron e 0 -1 Always outside the nucleus

Neutron n 1 0 Always in the nucleus

Proton p 1 +1 Always in the nucleus

Protons/Mass number/Isotope The atomic number of an element is the

number of protons in an atom of that element.

The mass number of an element is the sum of the number of protons & the number of neutrons in one atom of that element.

Isotopes of an element are atoms of that element containing the same number of protons but different numbers of neutrons.

Standard Atomic Notation

Radioactivity

If there are the “wrong” number of neutrons in a nucleus, the atom changes to get the number right. The atom fires out one or more fragments from its nucleus. This type of change is called radioactivity.

The changing radioactive atoms are said to

decay.

Electron shells

The electrons are not moving at random in the space around the nucleus.

They are arranged in layers, one outside the other. These layers of electrons are called electron shells.

A shell can only hold up to a maximum number of electrons. (2n^2)

1rst shell (2electrons); 2nd shell (8electrons) 3rd shell (18electrons); 4th shell (32

electrons)

Atoms & ions

Atoms are electrically neutral. Removing an electron from an atom

leaves the atom with a net positive charge.

Adding an extra electron to an atom gives it a net negative charge.

Charged atoms or groups of atoms are called ions.

Atoms vs. ions

Symbols for some common atoms & ions

Atom or ion

symbol Proton #

# of electrons

Net charge

Chlorine atom

Cl 17 17 0

Chloride ion Cl- 17 18 -1

Iron atom Fe 26 26 0

Iron ion Fe2+ 26 24 +2

Nitrate ion NO3- * * -1

Oxide ion O2- 8 10 -2

Sodium atom

Na 11 11 0

Sodium ion Na+ 11 10 +1

Sulphate ion SO42- * * -2

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 protons10 electrons

Cl 17 protons17 electrons

Cl-17 protons18 electrons

2.5

HISTORY OF THE ATOMHISTORY OF THE ATOM

460 BC Democritus develops the idea of atoms

he pounded up materials in his pestle and

mortar until he had reduced them to

smaller and smaller particles which he

called

ATOMAATOMA

(greek for indivisible)

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

particles called atoms. All atoms of a given element are identical. The atoms of one element are different from the atoms of all other elements.

2. Compounds are composed of atoms of more than one element. The relative number of atoms of each element in a given compound is always the same.

3. Chemical reactions only involve the rearrangement of atoms. Atoms are not created or destroyed in chemical reactions.

2.1

8 X2Y16 X 8 Y+

2.1

2

2.1

HISTORY OF THE ATOMHISTORY OF THE ATOM

1913 Niels Bohr

studied under Rutherford at the Victoria

University in Manchester.

Bohr refined Rutherford's idea by

adding that the electrons were in

orbits. Rather like planets orbiting the

sun. With each orbit only able to

contain a set number of electrons.

Bohr’s Atom

electrons in orbits

nucleus

HELIUM ATOM

+N

N

+-

-

proton

electron

neutron

Shell

What do these particles consist of?

ATOMIC STRUCTUREATOMIC STRUCTURE

Particle

proton

neutron

electron

Charge

+ ve charge

-ve charge

No charge

1

1

nil

Mass

ATOMIC STRUCTUREATOMIC STRUCTURE

HeHethe number of protons and neutrons in an atom

44Atomic mass number

the number of protons in an atom

22Atomic number

number of electrons = number of protons

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

2.2

Chadwick’s Experiment (1932)

H atoms - 1 p; He atoms - 2 p

mass He/mass H should = 2

measured mass He/mass H = 4

+ 9Be 1n + 12C + energy

neutron (n) is neutral (charge = 0)

n mass ~ p mass = 1.67 x 10-24 g

2.2

Subatomic Particles

Particle Mass

(g) Charge

(Coulombs) Charge (units)

Electron (e-) 9.1 x 10-28 -1.6 x 10-19 -1

Proton (p) 1.67 x 10-24 +1.6 x 10-19 +1

Neutron (n) 1.67 x 10-24 0 0

mass p = mass n = 1840 x mass e-

2.2

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 the nucleus

XAZ

H11 H (D)2

1 H (T)31

U23592 U238

92

Mass Number

Atomic NumberElement Symbol

2.3

2.3

Period

Group

Alkali M

etal

Noble G

as

Halogen

Alkali E

arth Metal

2.4

ATOMIC STRUCTUREATOMIC STRUCTURE

Electrons are arranged in Energy Levels

or Shells around the nucleus of an atom.

• first shell a maximum of 2 electrons

• second shell a maximum of 8

electrons

• third shell a maximum of 8

electrons

ATOMIC STRUCTUREATOMIC STRUCTURE

There are two ways to represent the atomic

structure of an element or compound;

1. Electronic Configuration

2. Dot & Cross Diagrams

ELECTRONIC CONFIGURATIONELECTRONIC CONFIGURATION

With electronic configuration elements are

represented numerically by the number of

electrons in their shells and number of shells. For

example;

N

Nitrogen

14

2 in 1st shell

5 in 2nd shell

configuration = 2 , 5

2 + 5 = 7

7

ELECTRONIC CONFIGURATIONELECTRONIC CONFIGURATION

Write the electronic configuration for the following elements;

Ca O

Cl Si

Na20

40

11

23

8

17

16

35

14

28B

11

5

a) b) c)

d) e) f)

2,8,8,2 2,8,1

2,8,7 2,8,4 2,3

2,6

DOT & CROSS DIAGRAMSDOT & CROSS DIAGRAMS

With Dot & Cross diagrams elements and

compounds are represented by Dots or Crosses to

show electrons, and circles to show the shells. For

example;

Nitrogen N XX X

X

XX

X

N7

14

DOT & CROSS DIAGRAMSDOT & CROSS DIAGRAMS

Draw the Dot & Cross diagrams for the following elements;

O Cl8 17

16 35a) b)

O

X

XX

X

X

X

X

X

Cl

X

X

X

X X

X

XX

X

X

X

X

X

XX

X

X

X

SUMMARYSUMMARY

1. The Atomic Number of an atom = number of

protons in the nucleus.

2. The Atomic Mass of an atom = number of

Protons + Neutrons in the nucleus.

3. The number of Protons = Number of Electrons.

4. Electrons orbit the nucleus in shells.

5. Each shell can only carry a set number of electrons.

2.6

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

2.6

ionic compounds consist of a cation and an anion

e.g. NaCl is really composed of Na+ & Cl _

• the formula is always the same as the empirical formula

• the sum of the charges on the cation and anion in each formula unit must equal zero… Is this always true?

The ionic compound NaCl

2.6

The term molecule is not normally used to apply to ionic substances.. Can you guess why?

The?

Formula of Ionic Compounds

Al2O3

2.6

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-

Some Polyatomic Ions

NH4+ ammonium SO4

2- sulfate

CO32- carbonate SO3

2- sulfite

HCO3- bicarbonate NO3

- nitrate

ClO3- chlorate NO2

- nitrite

Cr2O72-

dichromate SCN- thiocyanate

CrO42- chromate OH- hydroxide

2.7

Chemical Nomenclature• Ionic Compounds

– often a metal + nonmetal– anion (nonmetal), add “ide” to element name

BaCl2 barium chloride

K2O potassium oxide

Mg(OH)2magnesium hydroxide

KNO3 potassium nitrate

2.7

• 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

2.7

• Molecular compounds1. nonmetals or nonmetals + metalloids

2. common names

• H2O, NH3, CH4, C60

3. element further left in periodic table is 1st

4. element closest to bottom of group is 1st

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

6. last element ends in “ide”… this only applies for members of Group 5, 6, & 7

2.7

HI hydrogen iodide

NF3 nitrogen trifluoride

SO2 sulfur dioxide

N2Cl4 dinitrogen tetrachloride

NO2 nitrogen dioxide

N2O dinitrogen monoxide

Molecular Compounds

2.7

TOXIC!

Laughing Gas

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

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

2.5

Quality is never an accident. It is always the result of high intention, sincere effort and skillful execution. It represents the wise choice of many alternatives. 

unknown