IB 1 CHEMISTRY

Teacher: Annika Nybergannika.nyberg@mattliden.fi

Course book: Pearson Baccalaureate:Higher Level Chemistry for the IB Diploma2nd edition ISBN10: 1447959752

isbn13: 9781447959755

Pearson Baccalaureate:Standard Level Chemistry for the IB Diploma2nd editionISBN 10: 1447959760Isbn13: 9781447959762

http://www.chem1.com/acad/webtext/virtualtextbook.html

1. Stoichiometric Relationships

NOS: The atomic theory

● All matter is composed of atoms.

● These atoms cannot be created or destroyed during chemical reactions, they can only be rearranged.

● Physical and chemical properties of matter depend on the bonding and arrangement of these atoms.

NOS: The phlogiston theory

https://www.youtube.com/watch?v=3MuMPLoQZN4

The atom• Atoms are composed of subatomic

particles: protons, neutrons and electrons.

• Protons and neutrons are found in the nucleus of the atom and are therefore called nucleons.

In a neutral atom, the number of electronsequals the number of protons.

1.1 Matter

● The characteristics of matter:

➢ has a mass➢ occupies a volume in space➢ is made up of particles (atoms, moleules or

ions)➢ particles are in constant motion

States of matter

• Substances have different appearances and physical properties in different states of matter.

http://www.youtube.com/watch?v=s-KvoVzukHo&feature=related

Temperature

● The way the particles of matter move depends on the temperature.

temperature (K) = temperature (ºC) + 273,15

Changes of state

http://www.youtube.com/watch?v=pP_lZaOchE0Dry ice bubbles

http://www.youtube.com/watch?v=6s0b_keOiOUIce melting

Heating curve for water

The philosophers stonehttps://www.youtube.com/watch?v=zTsoClR49UI

Greatest discoveries chemistryhttps://www.youtube.com/watch?v=s7xxMX4Ovig

Elements and compounds

● An element contains atoms of only one type.

● Compounds are made up of more than one element, chemically bonded together.

● The properties of a compound is very different from those of its constituent elements, e.g. NaCl.

+ →

● Elements and compounds have a constant composition and are pure substances.

● Pure substances can combine to form a mixture.

Mixtures

• Mixtures can be either:

The language of chemistry

● Chemistry has a universal language.

● The International Union of Pure and Applied Chemistry (IUPAC) is an organization that develops a system of standardized nomenclature for chemical compounds.

● See the IUPAC gold book for chemical terminology: http://goldbook.iupac.org/index.html

氯化钠 سدیم کلرید Cloruro di sodio

http://fa.wikipedia.org/wiki/%D8%B3%D8%AF%DB%8C%D9%85_%DA%A9%D9%84%D8%B1%DB%8C%D8%AF

http://en.wikipedia.org/wiki/Sodium_chloride

Writing equations

● Write the correct formulas for all the reacting species, reactants on the left-hand side and the products on the right-hand side.

● Write the correct coefficients in front of each species. The reaction is then said to be stoichometrically balanced.

● It is good practice to include the state symbols: (s), (l), (g), (aq).

● Ions that remain unreacted in the reaction (= spectator ions) can be left out from the equation.

Ionic compounds

● In forming ionic compounds the number of ions used is such that the number of positive charges is equal to the number of negative charges → the compound is electrically neutral, e.g.

copper(II)carbonate

sodium sulfate

Molecules

● Some elements exist as molecules = one or more atoms held together by covalent bonds.

The atom economy

● In an ideal chemical process the amount of products equals the amount of reactants and no atoms are “wasted”.

percentage = Moleculas mass of useful products x 100%

atom economy Moleculas mass of atoms in reactants

https://www.youtube.com/watch?v=Zuyk4hfbjSA

1.2 The mole concept

● SI: the international system of measurment.

● The SI (Systeme International d´Unités) system has seven base units. All the other units are derived from them.

The gram was originally defined in 1795 as the mass of one cubic centimeter of water at 4 °C, making the kilogram equal to the mass of one liter of water. The prototype kilogram, manufactured in 1799 and from which the current kilogram is based, has a mass equal to the mass of 1.000025 liters of water.

● A single atom of an element has an extremely small mass, far too small to weigh.

● Ex. Calculate the mass of one hydrogen-1 atom.

1,67 ▪ 10-27 kg (+ 9,1 10-31 kg) = 0,00000000000000000000000167 g

● Chemists measure amount of substances in moles, by counting particles (atoms, molecules or ions).

Amount of a substance, n

● 12,00 grams of carbon-12 contains 6,02 ·1023 atoms. This number is called the Avogadro's constant (N

A or L).

● One mole of ANY substance contains 6,02 ·1023 particles.

Isotopes

● The atoms of an element are made up of an mixture of isotopes.

● Isotopes have: • - the same atomic number (= the same number of

protons and electrons)

• - different mass number (different number of neutrons)

• - similar chemical properties, but different physical properties (density, bp)

Relative atomic mass, Ar (no unit)

● The relative atomic mass of an element is the weighted mean mass of all the naturally occuring isotopes of that element relative to the mass of carbon-12.

● For example, hydrogen has 1/12 of the mass of carbon-12.

● 1 mol of hydrogen atoms has a mass of 1.01 g.

Sub-atomic particle Mass (kg) Relative mass

proton 1,67 ▪ 10-27 kg 1

neutron 1,67 ▪ 10-27 kg 1

electron 9,1 ▪ 10-31 kg 0,0005

Relative molecular mass, Mr

● The relative molecular mass, Mr, of a molecule

is the sum of the relative atomic masses of the atoms in the molecule (found in the periodic table).

● The relative formula mass, Mr, is similar to the

above but can be used with nonmolecular substances such as ionic compounds.

Mr (AgNO

3) =

Mr (C

2H

5OH) =

602 000 000 000 000 000 000 000

Molar masses, M (g mol-1)

● The mass of one mole of ANY substance is known as the molar mass.

● For example, 1 mol of iron contains 6.02 x 1023 iron atoms and has a mass of 55.85 g.

Formulas of compounds● The empirical formula is obtained experimentally by burning

a compound in oxygen so that all its elements forms oxides.

● The amount of oxides can be determined and that gives the original amount of each element.

● The empirical formula shows the simplest whole number ratio of atoms of each element in a particle of that substance, e.g. C

6H

12O

6

Formulas of compounds

● The empirical formula is obtained experimentally by burning a compound in oxygen so that all its elements forms oxides.

● The amount of oxides can be determined and that gives the original amount of each element.

● The empirical formula shows the simplest whole number ratio of atoms of each element in a particle of that substance, e.g. CH

2O

Molecular formula:

Structural formula:

● Remember to check that all the percentages add up to 100%!!

Ex. 17Calculate the empirical formula of a compound that contains 40.4% carbon, 6.0 % hydrogen, 17.7% nitrogen and 35.9% oxygen by mass.

● Ex.18A compound contains 12.79% carbon, 2.15% hydrogen and 85.06% bromine by mass. Its relative molecular mass is 187,9. Determine a) the empirical formula and b) the molecular formula of the compound.

g)

h)

i)

j)

k)

1.3 Reacting masses and volumes

● Chemical equations show reactants combining in fixed ratios (moles) to form products.

● Ex. Methane burns in air:

CH4 (g) + 2 O

2(g) → CO

2 (g) + 2H

2O (g)

● Calculate the mass of carbon dioxide produced from the complete combustion of 1,00 g of methane.

Theoretical yield

● The balanced chemical equation can be used to predict how much product can theoretically be produced from given masses of starting material.

Ex 7. a) Write the balanced chemical equation for the reaction where ethene and steam react to produce ethanol (C

2H

5OH).

b) What is the maximum amount of ethanol that can be produced when 1,0 kg of ethene and 0,010 kg of steam are placed into the reaction vessel?

Percentage yield

● Some liquids are pure substances, but more commonly liquids are solutions containing two or more components.

● Solution: homogenous mixture of two or more substances.

● solid/solid:● solid/liquid:● liquid/liquid:● gas/liquid:

1.5 Solutions

Concentration

● The concentration of a solution is the amount of solute (in moles) per volume of solution.

Dilution of solutions

● Stock solution: a concentrated starting solution

Density

density = _mass

volume

Acid-base titration

Ex. 12 Acid-base titrations:

● Sodium hydroxide reacts with hydrochloric acid according to the following equation:

NaOH (aq) + HCl (aq) → NaCl (aq) + H2O (l)

● Calculate the volume of 0,0500 mol dm-3 sodium hydroxide solution to react exactly with 25 cm3 of 0,20 mol dm-3 hydrochloric acid.

1.4 Mass and gaseous volume relationships in chemical

reactions

Properties of gases:

● Gases have a small mass

● All gases respond in a similar way to changes in temperature, pressure and volume.

● They exert a pressure, that depends on the amount of gas and the temperature

● There is no bonding between molecules

● The molecules may move in all directions allowing the gas to expand throughout any container

Pressure

● Pressure is the amount of force exerted on one unit of area.

Avogadro´s law

● One mole of any gas will occupy the

same volume, if the temperature

and pressure are the same.

= equal volumes of different gases at the same temperature and pressure contain the same number of particles.

● This volume is known as the molar volume of a gas, V

molar (unit dm3/mol)

n = V / Vmolar

● At standard temperature and pressure, STP (273 K and 101,3 kPa) one mole gas occupies 22,4 dm3.

● At room temperature, RTP (298 K and 101,3 kPa) one mole gas occupies 24 dm3.

● Ex. 16 s. 20: Calculate the volume occupied by 4,40 g of carbon dioxide at STP.

● Worked example, s.21: What volume of hydrogen is produced when 0,056 g of lithium reacts completely with water at STP:

2Li (s) + 2 H2O (l) 2LiOH (aq) + H

2 (g)

● Ex. 17 s. 21: Calcium reacts with water to produce hydrogen:

Ca (s) + 2H2O (l) Ca(OH)

2 (aq) + H

2 (g)

Calculate the volume of gas, measured at STP produced when 0,200 g of calcium reacts completely with water.

Ex 6.

Boyles law

● i.e. the relationship between volume and pressure for a gas

● The law describes how the volume of a given amount of gas at constant temperature varies inversely with the applied pressure:

PV = k1

http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/gasesv6.swf

Gay Lussac's law

● i.e. the relationship between temperature and pressure for a gas

● states that the pressure of a given amount of gas held at constant volume is directly proportional to its temperature in kelvin

= an increase in temperature increases the kinetic energy of the particles, which means they will move faster and collide with the walls with more energy and more frequency

P = k2T

Charles' law

● i.e. the effect of temperature on the gas volume

● The law describes how the volume of a given amount of gas is directly proportional to its temperature in kelvins.

V = k3T

The combined gas law

● The three gas laws can be combined to one expression:

where 1 refers to the initial conditions and 2 to the final conditions.

● Only ideal gases obey the gas laws perfectly. Realgases can be treated as ideal gases, unless dealingwith extremely precise measurements.

Ex 7.

The behaviour of gases

● There is large intermolecular distance between the molecules in a gas = a gas is mainly empty space.

● The gas molecules are free to move randomly in all directions. The molecules travel in a straight line until they collide with other gas molecules or with the walls of the container.

Kinetic theory of gases

● There are no attractive or repulsive forces between the atoms or molecules in a gas.

● The kinetic energy of a gaseous molecule or atom is given by the expression E = ½ mv2, where m is the mass of the particle and v is its speed.

● The mean kinetic energy of molecules or atoms in a gas is directly proportional to its temperature.

Ideal gas

● A gas that obeys all the gas laws under all conditions is said to be an ideal gas.

● No real gas behaves exactly as an ideal gas, but at high temperatures and low pressures the model describes most gases well.

●

● At high pressure and low temperature the gas particles are compressed:

- the gas particles move slowly and come so close together that intermolecular forces attract them to each other

- the actual volume occupied by the particles will be significant compared to the total volume of the gas.

The ideal gas equation● The ideal gas law relates pressure, volume,

temperature and amount of substance:

where R is the gas constant

= 8,31451 J K-1 mol-1

● Only ideal gases will follow

this equation exactly.

Ex. 8. At 273 K and 101 325 Pa, 12,64 grams of a gas occupy 4,00 dm3. Calculate the molar mass and relative molecular mass of the gas.

Ex 9.

Ex 10.

Ex 11.

● Aim 8: The negative environmental impacts of refridgeration and air conditioning systems are significant.

● The use of CFCs as refridgerants has been a major contributor to ozone depeltion.

https://www.youtube.com/watch?v=c5mfbm9oREk&list=PLgGh_MSicdaHo8OwhRxyC5I3YvbWAYyKL&index=5

https://www.youtube.com/watch?v=XLY8m-dXOxo