ChemistryChpter2&3

7/30/2019 ChemistryChpter2&3

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Chapter 2 : The Structure of The Atom

2.1 Matter

1) Matter = has mass & occupies space

= tiny, discrete particles atomsions

molecules2) Atom = smallest particles of element (take part in chemical reaction)3) Molecules = exist as particle

= contain 2 or more atoms

4) Ions = exist as particleCharged positively

negatively

= liquid state/molten state= when atom loses 1 or more electron, it becomes ( + ) ion,when atom gains 1 or more electron, it becomes ( - ) ion



2.1.2 Kinetic Theory of Matter

1) All matter made up of tiny particle continuously & random2) Attractive forces happen when particles close together

3) Movement = affected by heat4) If temperature increases, motion of particles increases.

The higher the temperature rises, the faster the motion of particles.5) Freedom of motion = affected by intermolecular forces.

The stronger the force, the slower the motion.

SOLIDS- strong forces of attraction

- not able to move freely- Only vibrate

- Cannot break free- Fixed shape & volume

LIQUIDS

- further apart- not arranged in regular order

- force attraction weaker- moves easily & attract 1 another-take the shape of container

GASES- further apart & not close together

- no attraction between particles- move faster & freely in zigzag- collide with 1 another & walls of thecontainer

Characteristics of the three states of matter

MatterCharacteristics Solid Liquid Gas

Density High Moderate Low

Compressibility None Almost impossible Possible

Shape Definite Take shape of container None

Volume Definite Definite Variable

Diffusion The slowest Moderately fast The fastest



2.1.3 Change in State Of Matter

1) Melting • solid heated , getting hotter & hotter• particles = vibrate more until they break away from 1 another

• happens at boiling point

•At B.P , solid liquid

2) Evaporation • liquid heated, particles moves faster as temperature rises• few particles at the surface have enough energy to escape as a gas

3) Boiling •more heat supplied to liquid, temp increases until reach B.P.• particles throughout the liquid have enough energy to break away

from liquid to form steam.• occurs at a constant temperature

4) Condensation • gas is cooled, particles lose energy•move slowly•when collide, they do not have enough energy to bounce away again• stay close together

5) Freezing • liquid cooled, particles slow down more•

liquid form a solid at constant temp• temperature (freezing point)

6) Sublimation • iodine is heated, does not melt

• get enough energy to break away completely from solid & form agas

• take place when solid CO2 ( dry ice ) is heated



Effects of Heat on the Kinetic Energy of Particles

Solid Liquid Gas

1) Strong forces attraction

2) Hold the atoms or

molecules together.

3) Vibrate

4) Cannot break free

1) The molecules are close

together

2) Have enough energy to

move about.

1) Molecules are far apart

2) Moving so quickly

3) Do not really attract one

another.

• Kinetic energy of particles related to the speed of particles in motion• Particle moving faster have more kinetic energy than a slower moving particle

• Heat is form of energy. When a substances is heated, it will receive heat energy.Temperature increases.

• Heat energy cause the particles to move faster

• Kinetic theory = molecules move more quickly when substances is heated

• On heat, molecules receive heat from source. Will increase kinetic energy ofmolecule.

• If kinetic energy increases, temperature increases

2.2.2 Proton Number



1 ) Proton number = no of protons

2 ) Atoms is neutral3 ) No. of electron in a neutral atom = no. of protons

2.2.3 Nucleon Number

1 ) Determine the mass of atom.

No. of neutrons = nucleon no. – no. of protons

= nucleon no. – proton number

2 ) Atom with a bigger nucleon no. will be heavier than atom with a smaller nucleon no.

that will be more lighter.

2.2.4 Symbol Of Element

2.3 Isotopes

1) Definition = Same no. of protons but different number of neutrons

2) Have same chemical properties because their electron are arranged in same way3) Will not have same mass because different no. of neutrons

2.4.1 Valence Electrons

1) Outermost shell = valance shell

2) Electrons in outermost shell = valance electron3) All element in same group = similar chemical properties.

= same number of valence electron

Chapter 3 : Chemical Formulae and Equations

Proton = electron

Nucleon number = no. of protons + no. ofneutrons



3.1 Relative Atomic Mass and Relative Molecular Mass

1. Relative atomic mass, Ar

is the atomic mass of an atom when compared to a standard atom

2. Standard atom:

Hydrogen scale: hydrogen is the lightest atom of all and the mass of one hydrogen

atom was assigned 1 unit.

Weakness of Hydrogen scale

not too many elements can react readily with hydrogen,

the reactive masses of some elements were not accurate,

hydrogen exists as a gas at room temperature and

has a number of isotopes with different masses.

Helium scale: the second lightest atom of all and the mass of one helium atom was

assigned 1 unit.

Weakness of Helium scale

Mass of 1 helium atom = 4 times the mass of a hydrogen atom

So, mass of 1 helium atom = 4 times 1/12 mass of a carbon atom

helium exists as a gas at room temperature & helium is an inert gas.

Oxygen scale: chose as the standard atom to compare the masses of atoms

Weakness of Oxygen scale

the existence of three isotopes of oxygen were discovered,

natural oxygen (containing all the three isotopes) as the standard (Chemist) and

used the isotopes oxygen-16 as the standard (Physicists).

Carbon scale: standard atom of comparison internationally a carbon-12 atom is 12times heavier than an atom of hydrogen, used as the reference standard in mass



spectrometers, exists as a solid at room temperature, most abundant carbon isotope,

happening about 98.89% andcarbon-12 is close to the agreement based on oxygen.

3. Relative molecular mass, Mr

substances is the average mass of a molecule (two or more atoms) of the

substances when compared 1/12 with of the mass of a carbon-12 atom.

4. Relative formula mass, Fr

is for ionic compound which is calculated by adding up therelative atomic

masses of all the atoms.

5. Example:

Relative atomic mass, Ar of helium = 4

Relative molecular mass, Mr of CO2 = 12 + 2(16) = 44

Relative formula mass, Fr of NaCl = 23 + 35.5 = 58.5

Relative formula mass, Na2CO3·10H2O = 2(23) + 12 + 3(16) + 10 [2(1) + 16] = 286

1. Avogadro constant / Avogadro’s number is 6.02 x 1023



2. Atomic substances - all the particles are atoms.

Example: zinc (Zn), sodium (Na), aluminium (Al) and all noble gases, argon (Ar),

helium (He) and neon (Ne).

RAM (Relative Atomic Mass) of Na = 23

3. Molecular substances

Covalent compounds – the particles are molecules.

Example: carbon dioxide (CO2), water (H2O) and non-metal elements, iodine (I2),

nitrogen (N2) and oxygen (O2).

RMM (Relative Molecular Mass) of I2 = 127 + 127 = 254

4. Ionic substances

Ionic compounds – the particles are ions.

Example: sodium chloride (NaCl), hydrochloric acid (HCl) & potassium iodide (KI).

RFM (Relative Formula Mass) of HCl = 1 + 35.5 = 36.5

5. Avogadro’s Law / Gas Law

states that equal volumes of all gases contain the same numberof molecules under the same temperature and pressure.

Example: equal volumes of molecular hydrogen and nitrogen would contain the

same number of molecules under the same temperature and pressure.

6. Volume of gas (dm3)

Number of moles of gas x Molar volume

7. Room temperature and pressure (r.t.p.)

24 dm3 mol-1 (25°C and 1 atm)

Example: What is the volume of 5.0 mol helium gas at s.t.p.?

Volume of gas = Number of moles x Molar gas volume

= 5.0 mol x 24 dm3 mol-1

= 120 dm3

8. Standard temperature and pressure (s.t.p.) 22.4 dm3 mol-1 (0°C and 1 atm)



Example: What is the volume of 5.0 mol helium gas at s.t.p.?Volume of gas = Number of moles x Molar gas volume

= 5.0 mol x 22.4 dm3 mol-1

= 112 dm3

9. Mass (g) Number of moles x Molar mass10. Number of particles

Number of moles x Avogadro constant

11. Volume (dm3)

Number of moles x Molar volume

Empirical and Molecular Formulae

1. simplest ratio of atoms of each element that present in the compound & molecular

formulae indicate:

the types of the elements the symbols of the elements and the ratio of atoms or

moles of atoms of each element in a compound.

2. Molecular formula = (empirical formula)n [n is a positive number]

Compound Molecular formula n Empirical formula

Carbon dioxide CO2 1 (CO2) = CO2

Ethane CH3 2 (CH3)2 = C2H6

Propene CH2 3 (CH2)3 = C3H6

Glucose CH2O 6 (CH2O)6 = C6H12O6

Quinine C10H12NO 2 C20H24N2O2

4. Cations are positively-charged ions.

Charge Cations Formula



+1 Ammonium ion NH4+

+1 * Copper(I) ion Cu+

+1 Hydrogen ion H+

+1 Lithium ion Li+

+1 * Nickel(I) ion Ni+

+1 Potassium ion K+

+1 Silver ion Ag+

+1 Sodium ion Na+

+2 Barium ion Ba2+

+2 Calcium ion Ca2+

+2 * Copper(II) ion Cu2+

+2 * Iron(II) ion Fe2+

+2 * Lead(II) ion Pb2+

+2 Magnesium ion Mg2+

+2 * Manganese(II) ion Mn2+

+2 * Tin(II) ion Sn2+

+2 Zinc ion Zn2+

+3 Aluminium ion Al3+

+3 * Iron(III) ion Fe3+

+4 * Lead(IV) ion Pb4+

+4 * Tin(IV) ion Sn4+

5. Anions are negatively-charged ions.

Charge Anions Formula



-1 Bromide ion Br-

-1 Chloride ion Cl-

-1 Chlorate(V) ion ClO3-

-1 Ethanoate ion CH3COO-

-1 Fluoride ion F-

-1 Hydride ion H-

-1 Hydroxide ion OH-

-1 Iodide ion I-

-1 Manganate(VII) ion MnO4-

-1 Nitrate ion NO3-

-1 Nitrite ion NO2-

-2 Oxide ion O2-

-2 Carbonate ion CO32-

-2 Chromate(VI) ion CrO42-

-2 Sulphide ion S2-

-2 Sulphate ion SO42-

-2 Sulphite ion SO32-

-2 Thiosulphate ion S2O32-

-3 Nitride ion N3-

-3 Phosphate ion PO43-

-3 Phosphite ion PO33-

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