Chemistry Module 1 PDF

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Notes on the preliminary module 1 part of chemistry...

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1Chemistry Module 1: Properties and Structure of Matter Inquiry Question 1: How do the properties of substances help us to classify and separate them? 

Explore homogeneous mixtures and heterogeneous mixtures through practical investigations - Using separation techniques based on physical properties - Calculating percentage composition by weight of component elements and/or compounds

Percentage composition by mass of element = Mass of element =

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Investigate the nomenclature of inorganic substances using IUPAC naming conventions

Nomenclature of inorganic substances Chemical substances can be subdivided into organic and inorganic substances. Organic compounds are those in which one or more atoms of carbon are covalently linked to atoms of other elements, most commonly, hydrogen, oxygen and nitrogen. Inorganic compounds are any substances in which two or more chemical elements (excluding carbon) are combined in nearly always definite proportions. The only carbon-containing compounds that are not classified as organic include of carbides (silicon carbide), carbonates (calcium carbonate) and cyanides (sodium cyanide). When a compound consists of atoms of two different non-metal elements: the element that is further to the left of the periodic table comes first in the name, the second non-metal has the end of its name changed (oxygen becomes oxide) and a prefix is used to indicate the number of atoms of each element that is present in a molecule of the compound (H2S is dihydrogen sulphide). When a compound consists of atoms of a metal and a non-metal: the name of the metal is used first, the end of the non-metal name is changed (chlorine becomes chloride), some transition metals can have more than one valency and polyatomic ions consist of two or more atoms bonded together.

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Classify the elements based on their properties and position in the periodic table through their: - Physical properties - Chemical properties

Properties Chemical properties are the properties that change the composition of an element or compound in a chemical change. Physical properties are things that can be observed without changing the composition of the element or compound and are used to observe and describe matter (solid, liquid and gas).

Examples of chemical properties: heat of combustion, flammability, reactivity, oxidation and a chemical’s reaction with water. Examples of physical properties: freezing, melting and boiling point, hardness, malleability, density and solubility. Metals All metals are solid under normal conditions, apart from Mercury, a liquid and gallium, a liquid on hot days. Physical properties of metal

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State- solid, excluding exceptions Lustre (reflect light off their surface) Malleable (able to be bent and hammered into shapes) Ductile (able to be stretched into wires) Hard, excluding sodium and potassium Valency of 1-3 electrons in their outermost shells Excellent conductors of heat and electricity as they have free electrons, lead is the poorest conductor of heat High melting and boiling points

Chemical properties of metal

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Electropositive character (low ionisation energy and typically lose electrons)

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Dissolve in water to form metal hydroxides

React with acids to from metal salts and water

Non-metals Physical Properties of non-metals

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State- solid or gas, only bromine exists as a liquid at room temperature

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Not malleable or ductile Brittle Poor conductors of electricity and heat Not lustre (do not reflect light) Low boiling and melting points Seven exist as diatomic molecules: H2(g), N2(g), O2(g), F2(g), Cl2(g), Br2(l), I2(s)

Chemical properties of non-metals

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Non-metals have the tendency to share or gain electrons with other atoms and are electronegative. When reacting with metals, they tend to gain electrons and become anions.

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Compounds composed of entirely non-metals are covalent substances and form acidic or neutral oxides

Metalloids Metalloids have properties between metals and non-metals. They are all solid in room temperature and can form alloys with other metals. Poorer conductor of heat and electricity than metals, physical properties tend to be metallic but chemical properties tend to be non-metallic.

Inquiry Question 2: Why are atoms of elements different from one another? 

Investigate the basic structure of stable and unstable isotopes by examining: - Their position in the periodic table

Isotopes are forms of an element with different number of neutrons in the atom and thus different atomic mass, however, have the same number of protons, and the same atomic number as they are the same element. When the atomic number is larger than 82 is radioactive, releasing alpha, beta and gamma decay to become lead which is the last stable element. The larger the atomic number, the more unstable it becomes.

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The distribution of electrons, protons and neutrons in the atom

The stability of isotopes is determined by the number of particles in the nucleus and the ratio of neutrons to protons. For light elements, the stable neutron to proton ratio is approximately 1:1, for heavy elements, the stable neuron to proton ratio is approximately 1.5:1. Isotopes are unstable when they neutron to proton ratio is not in the stable ratio. Unstable isotopes are called radioisotopes.

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Representation of the symbol, atomic number and mass number (nucleon number) Mass number is neutrons + protons (symbol A) Atomic number is number of protons (symbol Z)

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Model the atom’s discrete energy levels, including electronic configuration and spdf notation

Discrete energy levels Each shell can contain only a fixed number of electrons: the first able to hold 2 electrons, the second able to hold 8 electrons, the third able to hold 18 electrons and so on. The general formula is the nth shell can hold 2(n^2) electrons.

Electron configuration Electron configuration is the distribution of electrons of an atom or molecule in atomic or molecular orbitals and is used to describe the orbitals of an atom in its ground state. Electrons must be in one energy level or another, they cannot have energies that are intermediate between two levels. The electron configuration is determined by allocating electrons to the lowest energy levels, then the next energy levels when the energy levels are filled. Stable electron configurations are in noble gases as their energy levels are completely filled and therefore, these elements virtually have no tendency to undergo chemical reactions.

Spdf Notation There are four different types of orbitals (s, p, d and f) which have different shapes, and one orbital can hold a maximum of 2 electrons. However, the s, d and f orbitals have different sublevels and hold more electrons. The s sublevel contains one orbital, the p sublevel consists of three orbitals and the d sublevel consists of 5 orbitals. Electron spin- The Pauli Exclusion Principle states that there can be no more than two electrons in any given orbital and that these must have opposing spins . Hunds Rule states that every orbital in a subshell must first be filled with one electron with the same spin before an orbital is filled with a second orbital.

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Calculate the relative atomic mass from isotopic composition

Isotopes are atoms of the same element that have different numbers of neutrons in their nuclei, therefore, have the same atomic number but different mass number. The individual isotopes of each element have a relative isotopic mass. Relative isotopic mass is calculated through.

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Investigate energy levels in atoms and ions through collecting primary data from a flame test using different ionic solutions of metals examining spectral evidence for the Bohr model and introducing the Schrodinger model

Flame test

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Barium = green Calcium = orange/red Copper = blue/green Potassium = Lilac (purple) Sodium = Orange

Electrons orbit the nucleus in distinct energy levels. If energy is supplied to the atom, some electrons may absorb a quantum (specific amount) of extra energy and jump to a higher (outer) energy level (excited state). When the excited electron falls back to its original position (ground state), they emit the energy equal to the energy absorbed through visible light. The amount of energy emitted determines the wavelength and colour of the light. The emission radiation is recorded in the atomic emission spectra and is called atomic emission spectroscopy.

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Investigates the properties of unstable isotopes using natural and human-made radioisotopes as examples: types of radiation and types of balanced nuclear reactions

Inquiry Question 3: Are there patterns in the properties of elements? 

Demonstrate, explain and predict the relationships in the observable trends in the physical and chemical properties of elements in periods and groups of the periodic table, including but not limited to: - State of matter at room temperature - Electronic configuration and atomic radii - First ionisation energy and electronegativity

Ionisation energy is the energy required to remove an electron from an atom . First ionisation energy of an element is the energy required to remove an electron from an atom of an element in the gas state. The lower the ionisation energy, the easier it is to remove an electron. Ionisation energy is measured in kJ/mol. If more than one electron is removed, it is the 2 nd and 3rd ionisation energy and so on. The 2 nd ionisation energy is always greater than the first as more energy is needed to remove a negative electron due to it having to overcome a greater electrostatic attraction. Trends 1st ionisation energy decreases down groups of the periodic table as the activity of metals increase down the group. 1st ionisation energy increases along the periods of the periodic table as when moving across a period, the atomic number increases meaning that the number of protons increases, increasing the electrostatic attraction between the positive nucleus and the negative electrons.

Electronegativity of an element is a measure of the ability of an atom to attract electrons towards itself. The stronger the valence electrons of an atom are attracted to the nucleus of the atom, the greater the electronegativity.

Trends Electronegativity decreases down groups of the periodic table as the size of the atom increases, outer electrons are further from the nucleus so bonding electrons are not attracted so tightly. Electronegativity increases across a period as the atomic radius decreases, the valency shell is closer to the nucleus, the electrons being accepted from other atoms are closer to the nucleus and they are more strongly attracted. Fluorine, oxygen and nitrogen have the greatest electronegativity which is measured using the Pauling scale from 0.7-4.0. Reactivity with water

Inquiry Question 4: What binds atoms together in elements and compounds? 

Investigate the role of electronegativity in determining the ionic or covalent nature of bonds between atoms

When the difference in electronegativity is 1.5-4 it is an ionic bond, 0.3-1.5 it is a polar covalent bond and 0.0-0.3 is a non-polar covalent.

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Investigate the differences between ionic and covalent compounds through: Using nomenclature, valency and chemical formulae (including Lewis dot diagrams) - Examining the spectrum of bonds between atoms with varying degrees of polarity with respect to their constituent elements’ positions on the periodic table - Modelling the shape of molecular substances Investigate elements that posses the physical property of allotropy

Allotropes are forms of the one element (in the same physical state) that have distinctly different physical properties (colour, hardness, density, electrical conductivity).

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- Diamond and graphite are allotropes of carbon Investigate the different chemical structures of atoms and elements, including but not limited to: - Ionic substances High melting and boiling points- forces between the particles are strong All solid in room temperature Are hard- strong bonds Brittle- shatter easily Do not conduct electricity in the solid state- no free-moving charged particles Are good conductors of electricity in the liquid or molten state or when dissolved in waterfree-moving charged particles

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Covalent lattices (including diamond and silicon dioxide) Similar to ionic lattices in structure Extremel y hard Extremel y high melting and boiling points required to break the strong

covalent bonds holding the atoms together in the structure Do not conduct electricity or heat as there are no mobile charge carriers Very hard and brittle- very strong covalent bonds between the atom requires a large amount of energy to disrupt

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Low melting and boiling point Soft solids Generally gases Not conductive in any state or in solution Strength of intermolecular forces varies

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Covalent molecular

Metallic structure

All solid in room temperature, excluding Mercury Very good conductors of heat and electricity due to delocalised electrons Are malleable and ductile Most have high melting points and are fairly hard

Module 2: Introduction to Quantitative Chemistry Chemical Reactions and Stoichiometry IQ: What happens in chemical reactions? When a chemical reaction occurs, bonds are broken and bonds are made. Reactants  products. When a chemical reaction occurs, NEW products are formed. Stoichiometry: the study of quantitative aspects of formula and equations. 

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Conduct practical investigations to observe and measure the quantitative relationships of chemical reactions - Masses of solids and/or liquids in chemical reactions - Volumes of gases in chemical reactions Relate stoichiometry to the law of conservation of mass in chemical reactions by investigating:

The Law of Conservation of Mass: matter can be neither created nor destroyed, but merely changed from one form to another. In a chemical reaction, mass is conserved. Total mass of products = total mass of reactants. Conservation of mass requires that the total mass of the components of a chemical reaction in a closed system remains constant.

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Balancing chemical equations

As mass is conserved, the total number of each type of atom on the reactant side must be equal to the total number of each type of atom on the product side.

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Solving problems regarding mass changes in chemical reactions

Mole Concept 

Conducts a practical investigation to demonstrate and calculate the molar mass (mass of one mole) of: - An element - A compound

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Conduct an investigation to determine that chemicals react in simple whole number ratios by moles Explores the concept of the mole and relates this to Avogadro’s constant to describe, calculate and manipulate masses, chemical amounts and numbers of particles in:

Avogadro’s number ( N A ) is the number of particles in a mole (6.022 x 1023 )

n=

m MM -

Moles of elements and compounds

n=

m MM

where: n = number of moles, m = mass (g), MM = molar mass (g/mol) Percentage composition calculates and empirical formulae

The chemical formula of a compound represents the ration in which the atoms are present. The ratio in which elements are present is the percentage composition by weight.

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Limiting reagent reactions

Concentration and Molarity IQ: How are chemicals in solutions measured? 

Conduct practical investigations to determine the concentrations of solutions and investigate the different ways in which concentrations are measured

Solvent: substance that does the dissolving Solute: substance that is dissolved 

Manipulate variables and solve problems to calculate concentration, mass or volume using:

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c=

n V

Concentration equals n over v, (mol/L) Concentration also equals m over volume, (g/L) Different ways to measure concentration o

Percent by weight (%w/w)

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Percent by volume (%v/v)

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Parts per million, units of mg/L

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Dilutions (number of moles before dilution = number of moles of sample after dilution) Conduct an investigation to make a standard solution and perform a dilution

Gas Laws IQ: How does the Ideal Gas Law relate to all other Gas Laws? 

Conduct investigations and solve problems to determine the relationship between the Ideal Gas Law and:

Gas has a low density, a volume that fills the space available as particles move independently of one another, compress easily and has the ability to mix together rapidly.

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Gay-Lussac’s Law

In 1808, Gay-Lussac proposed the law known as Gay-Lussac’s law of combining volumes. This is when measured at constant temperature and pressure, the volume of gases taking part in a chemical reaction show simple whole number ratios to each other.

P =k . Pressure is proportional to T

Temperature.

P 2 P1 = T2 T 1

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Boyle’s Law

In 1662, Robert Boyle discovered by measuring the pressures and volumes of different gases at a constant temperature, the product of its Volume and Pressure is a constant. PV = k. No matter the gas, if there is a constant temperature and the amount of gas is not changed, the product of volume and pressure is a constant. Volume is inversely proportional to pressure, i.e. if pressure is doubled, volume is halved. P1 V 1=P 2 V 2

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Charles’ Law

In 1787, Jacques Charles discovered that a fixed quantity of gas at constant pressure, volume increases linearly with temperature-Charles Law. V=kT.

V =k . T

V2 V1 = T 2 T1 At constant pressure, the volume of a fixed quantity of gas is proportional to its absolute temperature.

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Avogadro’s Law

Equal volumes of gases contain equal numbers of molecules at the same temperature and pressure. A mole of gas has the same volume at constant temperature and pressure as a mole of any other gas. The volume of a gas is proportional to the number of moles present at a constant temperature and pressure independent of the nature of the gas. V=kn

When combining Boyle’s and Charles’s Laws, they give the combined gas law:

P1 V 1 P 2 V 2 = T2 T1 Ideal Gas Law

PV =nRT Constant R = 8.314J/K/mol

Separation techniques

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Distillation Evaporation Precipitation Filtration: consists of a filter funnel, filter paper that separate small particles from the solution, this flows into the beaker and is named the filtrate Chromatography: process of separating small amounts of substances from mixtures by the rates at which they move through or along a medium Magnetic: can be used to separate iron filings Centrifugation: done for solutions that need to be separated through density, mixture is spun fast and different, heavy parts of it sinks to the bottom and light, Vaporisation: followed by condensation, solution is heated to a vapor into a tube that is directed right into a glass apparatus that collects the steam and turns...