Unit 1 & 2 Chemistry Exam
calculate the percentage error when the experimental result can be compared with a theoretical or accepted result (value)
% error = theoretical - experimental /theoretical
apply the valence shell electron pair repulsion (VSEPR) theory to predict, draw and explain the shapes of molecules
yep
appreciate the law of conservation of mass and understand that the mole concept relates to mass, moles and molar mass
yep
calculate the measurement uncertainties in processed data, including the use of absolute uncertainties and percentage uncertainties
yep
communicate measurement uncertainties as a range (±) to an appropriate precision
yep
understand that hydrocarbons, including alkanes (saturated), alkenes (unsaturated) and benzene, have different chemical properties that are determined by the nature of the bonding within the molecules [chapter 7]
- Hydrocarbons are compounds composed entirely of carbon and hydrogen atoms. Because of the small electronegativity difference between carbon and hydrogen, all hydrocarbons are non-polar with dispersion forces acting between molecules. - Alkanes (saturated hydrocarbons) contain the maximum number of hydrogen atoms possible for the given number of carbon atoms. Alkanes are largely unreactive, which makes them good solvents for reactions involving otherwise highly reactive chemicals. - The main reaction of alkanes is the combustion, which is the reaction with oxygen to produce carbon dioxide and water and liberate heat energy. The energy required to break the C-C and C-H bonds is much less than the energy released from the production of C=O and O-H bonds. - Alkenes (unsaturated hydrocarbons) contain one or more double bonds between carbon atoms. - Benzene (C6H6) is composed of six carbon atoms arranged in a ring, with one hydrogen atom attached to each carbon atom. The extra stability that comes from these delocalised electrons explains why benzene does not readily undergo the addition reactions that are usually seen for alkenes. These substitution reactions often require quite reactive reagents or higher temperatures than are required for substitution reactions of alkenes; this is because of the high stability of benzene.
analyse and interpret given data to evaluate the properties, structure and bonding of ionic, covalent and metallic compounds [chapter 7]
- Ionic compounds are held together as solids by the attractions between oppositely charged ions. The ions are held in a repeating three-dimensional arrangement called a crystalline lattice structure. - Ionic compounds do not conduct electricity in the solid-state but can conduct electricity if dissolved in solution or in the molten state. - Ionic compounds usually have relatively high melting and boiling points, reflecting the strength of the electrostatic interactions between ions. Ionic compounds as solids are relatively hard but brittle. Many ionic compounds are water-soluble. - Metals in the free electron model of metallic bonding are regarded as positively charged ions surrounded by a sea of delocalised electrons. The electrostatic attractions between the cations and negatively charged electrons hold metallic solids together. Metals are malleable, ductile and good conductors of electricity. Generally, they have high boiling points. - The boiling points of metals decrease going down the first two groups of the periodic table. This is because as the atomic radius increases, the electrostatic attraction decreases. - Covalent bonds occur when non-metal atoms share electrons. Simple molecules with covalent bonds have low melting and boiling points and are non-conductive. - Giant covalent networks consist of many covalent bonds between atoms and generally, cannot conduct electricity (although there are exceptions such as graphite). They also have high melting points.
use data to calculate the enthalpy change (ΔH) for a reaction
- energy of products (-) energy of reactant
understand what mass spectrometry involves and that how the spectra generated can be analysed to determine and interpret certain things [chapter 5]
- mass spectrometry: an analytical technique in which chemical elements are ionised and the ions are sorted according to their mass-to-charge ratio - involves the ionisation of substances and the separation and detection of the resulting ions - a sample of an element is injected into the ionisation chamber of a mass spectrometer: 1. The element is bombarded with electrons, leaving the element positively charged, having lost an electron and formed an ion (ionisation) 2. The ions then pass through an electric field, which accelerates them through a magnetic field 3. The ions follow a curved path, with the paths of more positively charged or lighter ions bending more than the paths of less positively charged, heavier ions - spectra generated can be analysed to determine the isotopic composition of elements and interpreted to determine the relative atomic mass (analysis to determine)
understand how chemical bonds are caused by and what valency is [chapter 3]
A chemical bond is a lasting attraction between atoms, ions or molecules that enables the formation of chemical compounds. The bond may result from the electrostatic force of attraction between oppositely charged ions as in ionic bonds or through the sharing of electrons as in covalent bonds - chemical bonds are formed by electrostatic attractions that arise because of the sharing or transfer of electrons between participating atoms (allows atoms to become stable Valency refers to the ability of an atom or a group of chemically bonded atoms to form chemical bonds with other atoms or groups of atoms - the valency is a measure of the number of bonds that an atom can form (the number of electrons which can fill its valence electron shell)
analyse, interpret and evaluate data from flame tests and atomic absorption spectroscopy (AAS) to determine the presence and concentration of metallic ions in solution [chapter 5]
AAS is a quantitative technique because it can measure the concentration of metal atoms in a sample. The sample is compared to a set of standards of known concentration of the metal being analysed in the sample. A calibration curve of the standards is prepared, and the unknown sample's concentration is calculated from this. The calibration curve is a graph that plots the concentration of the standards on the x-axis and the absorbance of the light on the y-axis.
understand what flame tests and atomic absorption spectroscopy (AAS) are used for [chapter 5]
Absorption and emission spectra are essential to the qualitative and quantitative analysis of metals. The flame test and atomic absorption spectroscopy (AAS) are analytical techniques that can be used to identify elements. They both rely on electron transfer between atomic energy levels and are shown by line spectra. flame test: analyses metal ions (cations) by heating them in a bunsen burner flame and measuring the light emitted when excited electrons return to their ground state. Used to demonstrate the characteristic colours that metals emit when placed in a blue Bunsen burner flame, the heat of the flame excites electrons to a higher electron orbital shell and photons of light are emitted at distinct wavelengths when electrons return to their ground state. AAS: determines the concentration of metal atoms in the gas phase by measuring the amount of light absorbed from a specific lamp (absorption spectra of metals are used to determine the quantity of metals ions in an unknown sample). It works by spraying a solution of the sample into a flame, lamp is then set to a specific wavelength and the light beam excites the electrons in the atom. As electrons return to their ground state, they emit photons of light in which the detector measures the intensity (absorbance) of the light that is emitted, the greater the intensity, the higher the concentration of the metal atoms in the sample. - a set of standards with known concentrations are analysed and a calibration curve is created and the unknown sample is calculated from a graph
explain how varying the conditions present during chemical reactions, including temperature, surface area, pressure (gaseous systems), concentration and the presence of a catalyst can affect the rate of the reaction
An increase in concentration, increase in pressure (and therefore decrease in volume), increase in surface area or increase in temperature all result in a faster chemical reaction.
use the collision theory to explain and predict the effect of concentration, temperature, pressure and surface area on the rate of chemical reactions by considering the structure of the reactants and the energy of particles
An increase in concentration, increase in pressure (and therefore decrease in volume), increase in surface area or increase in temperature all result in a faster chemical reaction.
analyse, evaluate and interpret data to explain and justify conclusions for periodic trends, patterns and relationships [chapter 2]
Atomic radii and metallic character decrease across the period and increase down the group. Valency increases across the period and stays consistent down the group. Ionic radii are smaller than atomic radii when the ion is positive, and larger than atomic radii when the ion is negative. Ionisation energy and electronegativity increase across the period and decrease down the group. Successive ionisation energies refer to the removal of the first valence electron, then second, third and so on.
use and apply the nuclear symbol notation to determine the number of protons, neutrons and electrons in atoms, ions and isotopes [chapter 2]
Atoms - normal Isotopes - the same atom with a different mass number as there are a different number of neutrons Ions: - cations (metals that lose electrons) - anions (non-metals which gain electrons) The nuclear symbol notation can be used to determine the number of protons, neutrons and electrons in atoms, ions and isotopes.
construct and explain Maxwell-Boltzmann distribution curves for reactions with and without catalysts
Catalysts provide an alternative pathway, which requires less activation energy, resulting in an increased rate of chemical reaction. Catalysts can be represented on the Maxwell-Boltzmann distribution as well as on an energy profile diagram.
sketch and use energy profile diagrams, including the transitional state and catalysed and uncatalysed pathways, to represent the enthalpy changes and activation energy associated with a chemical reaction
Catalysts provide an alternative pathway, which requires less activation energy, resulting in an increased rate of chemical reaction. Catalysts can be represented on the Maxwell-Boltzmann distribution as well as on an energy profile diagram.
recognise that chromatography techniques, including paper, thin layer, gas and high-performance liquid chromatography, can be used to determine the composition and purity of substances
Chromatography is an analytical technique that can be used for qualitative or quantitative analysis. It can both identify components in a mixture and determine how much of each component is present. paper chromatography - an analytical technique for separating and identifying mixtures; the stationary phase is a thin strip of absorbent paper thin-layer chromatography (TLC) - an analytical technique for separating and identifying mixtures; the stationary phase is typically a thin layer of silica gel, aluminium oxide or cellulose supported on a piece of glass or plastic
explain how endothermic and exothermic reactions relate to the law of conservation of energy and the breaking and reforming of bonds; understand that heat energy is released or absorbed by the system to or from the surrounds
Endothermic Reaction: Energy is required Heat energy is absorbed from the surroundings, therefore feels cold Heat is found in the reactant Exothermic Reaction: Energy is released Heat energy is released into the surroundings, therefore feels hot Heat is found in the product Conservation of Energy: law of conservation of energy a law that states energy cannot be created or destroyed Therefore; reactions are just rearrangements of bonds
explain, in terms of average bond enthalpies, why reactions are exothermic or endothermic
Endothermic Reaction: Energy is required Heat energy is absorbed from the surroundings, therefore feels cold Heat is found in the reactant Ectothermic Reaction: Energy is released Heat energy is released into the surroundings, therefore feels hot Heat is found in the product
apply the relationship between temperature and enthalpy changes to identify thermochemical reactions as exothermic or endothermic; deduce from enthalpy level diagrams and thermochemical equations the relative stabilities of reactants and products, and the sign of the enthalpy change (ΔH) for a reaction
Enthalpy: chemical substances, referred to as its chemical energy or heat content Bond enthalpy: the amount of energy stored within a single bond ENDO → Enthalpy = +H ECTO → Enthalpy = -H Eact=Activation energy - energy required to make a reaction
understand that heat is a form of energy and that temperature is a measure of the average kinetic energy of the particles
Enthalpy: chemical substances, referred to as its chemical energy or heat content Bond enthalpy: the amount of energy stored within a single bond ENDO → Enthalpy = +H ECTO → Enthalpy = -H Eact=Activation energy - energy required to make a reaction
explain how catalysts, including enzymes and metal nanoparticles, affect the rate of certain reactions by providing an alternative reaction pathway with a reduced activation energy, hence increasing the proportion of collisions that lead to a chemical change
Enzymes and metal nanoparticles are examples of catalysts. Chemical reactions can be both fast and slow. Enzymes act as biological catalysts in a process that can be explained either by the lock model and key or induced-fit model.
distinguish between heterogeneous and homogeneous mixtures
Heterogeneous mixture = mixture with an inconsistent composition Homogeneous mixture = mixture with a uniform composition
understand that the type of bonding within ionic, metallic and covalent substances explains their physical properties, including melting and boiling point, thermal and electrical conductivity, strength and hardness [chapter 7]
Ionic bond: are held together by the electrostatic attraction between the oppositely charged ions (transfer electrons from metal cation - to non-metal anion) - strong bond = very high melting/boiling point, brittle (shatter when struct), a solid - good conductor of electricity, liquid - good conductivity Covalent bond: sharing electrons (non-metals) - covalent molecular (not joined together - weak intermolecular forces - easy to separate) = low melting point, brittle, no conductivity (no ions or free electrons) Metallic bond: holds metals together (free electrons shared by a lattice of cations - closely packed cations floating in a 'sea of electrons') - electrons not bound to atoms = high melting/boiling point due to strong attractions; conducts electricity and thermal heat; hard but malleable and ductile)
understand that the number of electrons lost, gained or shared is determined by the electron configuration of the atom and recall that transitional elements can form more than one ion [chapter 3]
Ionic bonding is the process by which atoms transfer valence electrons to each other. The transfer takes place between an atom ready to lose one or more electrons, and an atom ready to gain one or more electrons. Covalent bonding occurs when atoms share electrons to form bonds. Covalent bonds typically form between two or more non-metals and are influenced by the electronegativity values of each of the atoms involved.
explain how successive ionisation energy data is related to the electron configuration of an atom [chapter 2]
Ionisation energy (IE) is the energy required to remove an electron from a gaseous atom, forming a positively charged ion. Successive ionisation energies refer to the removal of the first valence electron, then second, third and so on. An atom's first IE is the energy required to remove one electron from its valence energy level. An atom's second IE is the energy required to remove a second electron from its valence energy level. Therefore, the valency of an atom has a significant impact on the differences between an atom's successive ionisation energies, depending on when an atom has a whole electron shell removed.
recall what isotopes are and how they can be represented [chapter 4]
Isotopes are atoms of an element that have the same number of protons but different numbers of neutrons. As a result, all isotopes of an element have the same atomic number but different mass numbers. AX or X-A: A - represents the mass number X - represents the elemental symbol (or elemental name) used to identify the number of protons contained in the atom.
recognise that isotopes of an element have similar chemical properties but have different physical properties [chapter 4]
Isotopes of the same element have the same electron configuration when neutrally charged. For this reason, the chemical properties of isotopes of an element are similar. However, physical properties are determined by the mass number. The varying number of neutrons causes isotopes to have different physical properties eg. heavy water D20 and light water H20
deduce Lewis (electron dot) structure of molecules and ions showing all valence electrons for up to four electron pairs for each atom [chapter 3]
Lewis structures (or electron dot structures) are visual diagrams that represent the valence electrons present in atoms, ions and molecules. 1. Count the total number of electrons in the molecule, from each of the atoms present. 2. Place the atoms' element symbols together, starting in the centre with the atom that has the most bonding electrons. Surround the central atom with the other atoms. The least electronegative atom is usually the central atom. 3. Draw a single line to connect the central atom with each of the surrounding atoms, to represent bonds. Draw two-electron dots on each of these bonding lines. 4. Draw the lone pairs of electrons on the surrounding atoms to complete the octet. 5. Draw any remaining electrons on the central atom. If there are not enough electrons to finish the central atom's octet, then multiple bonds need to be formed: double or triple bonds.
use appropriate mathematical representations to make inferences and to solve problems, including calculating the relative atomic mass of an element and percentage abundances of the isotopes of an element from data [chapter 5]
Mass spectrometry causes ions to be separated according to their mass-to-charge ratio (mass spectrum), mass spectrum displays the mass-to-charge ratio of each particle on the x-axis and its percentage abundance on the y-axis. The spectrum generated shows: - how many isotopes an element has because each has a different mass-to-charge ratio - the relative mass of each isotope, indicated by the mass-to-charge ratio - how much of each isotope is in the sample, indicated by the percentage abundance.
compare and explain the metallic and non-metallic behaviours of elements, including group trends and the reactivity for the alkali metals (Li-Cs) and the halogens (F-I) [chapter 2]
Metallic and non-metallic behaviour: - the more an atom wants to give away electrons, the more metallic it is; the more it wants to take an electron the more non-metallic it is The periodic table is divided into metals, metalloids and non-metals, as well as horizontal periods and vertical groups group 1, are the alkali metals, group 2 is the alkaline earth metals, groups 3-12 are the transition metals, group 17 is the halogens, and group 18 is the noble gases. Alkali metals: Silver coloured, soft metals with a low density. They easily react with water and halogens. Their reactivity increases down the group (Li-Cs), with caesium being the most reactive. Alkaline earth metals: Silver coloured, hard metals with a slightly higher density than alkali metals. They react slowly with water, sometimes requiring hot water or even steam. Halogens: Exist as non-metallic diatomic molecules in their natural state. They are highly reactive with alkali metals and can be harmful or lethal to the environment. The halogens decrease in reactivity down the group (F-At) with fluorine being the most reactive. Noble gases: mostly unreactive
understand what the relative atomic mass of an element is and how it is derived from the naturally occurring form of the element to 1/12 the mass of an atom of carbon-12 [chapter 4]
Most elements have two or more naturally occurring isotopes, but the atomic mass/mass number included in the periodic table is based on an average of the atomic masses of all the naturally occurring isotopes for an element. This number is called relative atomic mass (RAM) because it has been quantified relative to one-twelfth the atomic mass of the carbon-12 isotope. Calculating the relative atomic mass of an element uses the atomic masses of all of the naturally occurring isotopes of the element and the natural abundance (usually represented as a percentage abundance) of each of the isotopes.
analyse the impact of random error/measurement uncertainties and systematic errors in experimental work and evaluate how these errors/measurement uncertainties can be reduced
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recognise that oxides change from basic through amphoteric to acidic across a period [chapter 2]
Oxides are compounds that contain oxygen and another element, which can be a metal or non-metal. The element's degree of metallic or non-metallic properties determines what type of oxide the compound will be. As metallic character decreases across a period, oxides change from basic through amphoteric to acidic. basic - has a pH of higher than 7 (Basic oxides form between oxygen and metals) amphoteric - can be basic or acidic (Amphoteric oxides have both acidic and basic properties. They are metals that can form more than one oxidation state (or ionic charge) acidic - has a pH of lower than 7 (Acidic oxides form between oxygen and non-metals)
understand that the unique properties of water, including boiling point, density in solid and liquid phases, surface tension and ability to act as a solvent can be explained by its molecular shape and hydrogen bonding between molecules
Oxygen is electronegative, so it pulls on the electrons in the covalent bonds, meaning that they spend more time closer to the oxygen nucleus. This means that in a O-H bond, the negative charge is closer to the oxygen atom than the hydrogen atom. Because water is bent, a dipole is formed in which oxygen has a partial negative charge and each hydrogen has a partial positive charge. The presence of partial charges in a molecule makes it a polar molecule.
describe and explain that elements of the periodic table show trends across periods and down groups, including atomic radii, valencies, ionic radii, 1st ionisation energy and electronegativities as exemplified by groups 1, 2, 13-18 and period 3 [chapter 2]
Periodic trends are relationships between groups of elements and their properties, when structured in increasing atomic number and organised into periods and groups. 1. Atomic radius: refers to the size of an atom *measured in pm (picometres x10^-12m) → As Z increases, the atomic radius increases → Atomic radius decreases as you go left to right on the table → Atomic radius increases as you go down the table because of the electron shielding effect. 2. Valency: is the number of bonds an atom can make → Valency increases across a period, noble gases in any period having complete sublevels →Valency also depends on the group number 3. Ionisation energy: energy required to remove an electron from a gaseous atom to form a cation → if the radius decreases, electrostatic attraction between the nucleus + valence electrons increases. Therefore ionisation energy increases across the period. → Decreases as you go down the group, due to electron shielding effect. 4. Electronegativity: the attraction between (+ charged) nucleus and electrons in a neighbourly atom → electronegativity increases across a period → electronegativity decreases down a group
consider the relationship between the volume, number of moles and molar volume at standard temperature and pressure (STP)
STP - n = v/22.7
distinguish between the terms solute, solvent, solution, and concentration
Solute = a minor component in a solution in which dissolves into a solvent. Solvent = component which is able to dissolve other substances. Solution = homogeneous mixture composed of two or more substances where a mixture of a solute is dissolved in a solvent. Concentration = amount of a substance in a defined space divided by the total volume of a mixture.
apply the Aufbau principle, Hund's rule and the Pauli exclusion principle to write electron configurations for atoms and ions up to Z = 36 and use orbital diagrams to represent the character and relative energy of orbitals [chapter 2]
The Aufbau principle, Hund's rule and the Pauli exclusion principle are applied when writing electron configurations and creating orbital diagrams. Aufbau principle: electrons will fill the orbital with the lowest energy level Hund's rule: if two or more orbitals in the same sub-shell are available, one electron will go into each orbital (electron repulsion) Pauli exclusion: no two electrons can occupy the same space with the same spin (each orbital contains two electrons - one of each spin)
explain how emission spectra are produced [chapter 5]
The electromagnetic spectrum consists of the full range of frequencies of light. It displays the range of electromagnetic radiation, its energies, frequencies and wavelengths. Different wavelengths are refracted by different amounts of light. When an atom absorbs energy in the form of electromagnetic radiation, electrons move from a lower energy electron shell to a higher energy electron shell: this is called the excited state because electrons contain more energy. The more energy that is absorbed, the electrons travel further away from the nucleus. Emission spectra are produced when the light that the sample emits when the electrons move from an excited state to a lower state is focused through a slit and passed through a prism to disperse the light into a spectrum.
recognise that the properties of atoms, including their ability to form chemical bonds, are explained by the arrangement of electrons in the atom and by the stability of the valence electron shell [chapter 3]
The electron configuration of an atom can be separated into two basic components: the inner electrons and the valence (outer) electrons. Many of an element's properties result from the number of valence electrons in the element. For example, the reactivity of an element is determined by its electrostatic attraction to adjacent atoms. Its valence electrons are gained, lost or shared with these adjacent atoms in order to acquire a stable and low-energy noble gas electron configuration. The more stable the valence electron shell, the less likely the atom will want to react with other atoms. octet rule: atoms gain and lose electrons to reach eight electrons in their valence shells.
identify the numbers of bonding and lone pairs of electrons around each atom in a molecule [chapter 3]
The electrons in Lewis structures of covalent molecules are categorised as either singular bonding electrons or lone pairs of (non-bonding) electrons. bonding electrons - valence electrons that are bonded to another atom lone pair of electrons - valence electrons that are not bonded to another atom double bond - a bond in which two atoms share four electrons triple bond - a bond in which two atoms share six electrons A coordinate covalent bond is formed when the two electrons in the covalent bond have been sourced from only one of the atoms involved (represented by an arrow instead of a standard bond line).
analyse, interpret and evaluate data from chromatographs to determine the composition and purity of substances, including calculating Rf values
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distinguish between unsaturated, saturated and supersaturated solutions
Unsaturated solution = a solution that contains less than the maximum amount of solute that is capable of being dissolved. Saturated solution = a solution in which no more solute can be dissolved in the solvent. Supersaturated solution = that contains more of the dissolved material than could be dissolved by the solvent under normal circumstances. Has to be heated to dissolve more solute.
determine the formula of an ionic compound from the charges on the relative ions and name the compound [chapter 3]
When representing an ionic compound as an ionic formula, write the cation first, followed by the anion. Use the swap and drop method.
use the appropriate stoichiometric ratio to determine that reactants can be limiting
a limiting reagent is a substance that is totally consumed during a chemical reaction, therefore the amount of product is limited by the moles of the limiting reagent
understand and apply the reactions of acids with bases, metals and carbonates to determine reactants and products
acid - base = salt and water acid - metal = salt and hydrogen acid - carbonate = salt + water + carbon dioxide
use the pH scale to compare the levels of acidity or alkalinity of aqueous solutions
acidic = <ph 7 alkaline = >ph 7
understand how atoms can be modelled; what electron configurations show; and what the structure of the periodic table is based on [chapter 2]
atoms can be modelled as a nucleus surrounded by electrons in distinct energy levels (and sublevels) held together by electrostatic forces of attraction between the nucleus and electrons; inside the nucleus is the positively charged protons and neutral neutrons. The location of electrons within atoms can be represented using electron configurations, and the structure of the periodic table is based on the electronic configuration of atoms
recall that concentration can be represented in a variety of ways including the number of moles of the solute per litre of solution (mol L-1 ) and the mass of the solute per litre of solution (g L-1 ) or parts per million (ppm)
c = n/v
use appropriate mathematical representations to solve and make predictions (including using the mole concept and the relationship between the number of moles of solute, concentration and volume of a solution) to calculate unknown values
c=m/v
use appropriate mathematical representation to solve problems and make predictions, including determining the percentage composition from relative atomic masses; empirical formula of a compound from the percentage composition by mass; and molecular formula of a compound from its empirical formula and molar mass
determining the percentage composition from relative atomic masses the empirical formula of a compound from the percentage composition by mass the molecular formula of a compound from its empirical formula and molar mass
distinguish between absorption and emission spectra and recognise what the Bohr model provides evidence for [chapter 5]
electrons can absorb and emit electromagnetic radiation in the visible spectrum (prism is used to separate light into different wavelengths) absorption spectra - formed when white light passes through a sample and it absorbs specific wavelengths - this will leave gaps - viewed as a set of black lines emission spectra - formed when a sample is heated and the electrons go from an excited state back to ground state and certain wavelengths of light are produced - viewed as a set of coloured lines Bohr model: The absorption and emission spectra of the element hydrogen provide the evidence for Niels Bohr's theory that electrons orbit the nucleus of an atom in distinct shells. - electrons orbit an atom's nucleus - electrons exist in orbitals, which are distinct distances from the nucleus and have definite energy levels - for an electron to move between orbitals, it must gain energy or lose energy by absorbing or emitting photons of light with distinct wavelengths.
recall how elements are represented by symbols and recognise how the structure of the periodic table is based on them [chapter 2]
elements are represented by symbols and recognise that the structure of the periodic table is based on the atomic number and the properties of the elements Elements are represented by chemical symbols and are located in the periodic table according to the atomic number and properties of the elements. The periodic table is divided into metals, metalloids and non-metals, as well as horizontal periods and vertical groups group 1, are the alkali metals, group 2 is the alkaline earth metals, groups 3-12 are the transition metals, group 17 is the halogens, and group 18 is the noble gases.
recognise that a mole is a precisely defined quantity of matter equal to Avogadro's number of particles
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recognise the electron configuration of Cr and Cu as exceptions [chapter 2]
expected: Cr - 1s2 2s2 2p6 3s2 3p6 4s2 3d9 actual: Cr - 1s2 2s2 2p6 3s2 3p6 4s1 3d10 expected: Cr - 1s2 2s2 2p6 3s2 3p6 4s2 3d4 actual: Cu - 1s2 2s2 2p6 3s2 3p6 4s1 3d5 The expected electron configuration has 4s full and 3d partially full, whereas the actual electron configuration has both 4s and 3d half-full and therefore is more stable overall. This is because some energy is required to pair up two electrons in the same orbital and the 4s and 3d orbital are very close in energy. The same situation applies to copper: 4s becomes half-full so that 3d can become completely full.
appreciate that experimental yield can be different from theoretical yield
experiment/perecnt yield = actual amount of product/amount you would expect to get from the reaction
calculate percentage yield from experimental or given data
experiment/perecnt yield = actual amount of product/amount you would expect to get from the reaction
calculate the heat change for a substance given the mass, specific heat capacity and temperature change
heat change (q) - mcAt - mass of water x 4.18 x change in temp /1000 - kj - divide that by moles to get delta h
use the kinetic theory of gases to describe and explain the behaviour of gases, including the qualitative relationships between pressure, temperature and volume
ideal gas law: PV = nRt - pressure increase = temperature increase, volume decrease - temperature increase - pressure increases, volume increases - volume increase - pressure decrease, temperature increase
recognise that the properties of ionic compounds can be explained by modelling ionic bonding as ions arranged in a crystalline lattice [chapter 7]
ionic compound - a compound consisting of positively charged cations and negatively charged anions An ionic compound is an ordered, repeating three-dimensional arrangement of cations and anions in a crystalline lattice structure. In a crystalline lattice, cations are surrounded by anions and vice versa. This arrangement maximises the attraction between the oppositely charged ions. (metallic lattice, giant covalent networks, allotropes —carbon) ionic bonding - the electrostatic (charge-based) attractive forces holding ions together in ionic compounds. characteristics of ionic compounds: 1. Ionic compounds have high melting and boiling points as the ions are locked in place by the electrostatic interaction, so it takes a large amount of heat energy to break these bonds 2. The solids are relatively hard but brittle as under pressure, ions of the same charge can be brought close together, resulting in repulsion 3. Ionic compounds can conduct electricity as liquids as ions are no longer held in a rigid lattice and are free to move between the electrodes, allowing to carry the current. Aqueous solutions of ionic solids can also conduct electricity. When dissolved in water, the ions separate and move toward the electrode with the opposite charge.
recognise what ions are and recognise that ions are represented by formulas which include the number of constituent atoms and the charge of the ion [chapter 3]
ions are atoms or groups of atoms that are electrically charged due to an imbalance in the number of electrons and protons cation: a metal atom that has lost one or more electrons and is positively charged eg. sodium loses one electron, it develops a +1 ionic charge anion: a non-metal atom which has gained one or more electrons and has a negative charge eg. chlorine gains an electron it develops a -1 ionic charge monatomic ion - an ion with only one atom polyatomic ion - an ion with two or more atoms Use the groups in the periodic table to remember common ionic charges. Group 1 is +1, group 2 is +2, group 13 is +3, group 15 is -3, group 16 is -2 and group 17 is -1.
analyse and interpret given data to evaluate the physical properties of pure substances and mixtures
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analyse experimental data, including constructing and using appropriate graphical representations of relative changes in the concentration, volume and mass against timey
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recognise that materials are either pure substances with distinct measurable properties (e.g. melting and boiling point, reactivity, strength, density) or mixtures with properties dependent on the identity and relative amounts of the substances that make up the mixture
mixture - a substance that is made up of two or more substances that are not chemically bonded Pure substances have distinct, measurable properties. The purity of a substance can be established by collecting or analysing data of its physical and chemical properties, such as melting point, boiling point, and density. Pure substances have characteristic reactivities and bond strengths.
use appropriate mathematical representation to solve problems and make predictions, including the mole concept, to calculate the mass of chemicals and/or the volume of a gas (at standard temperature and pressure) involved in a chemical reaction
n = v/22.7
recall that pH is dependent on the concentration of hydrogen ions in solution
pH = -log[H+] [H+] = 10^-ph
recall that pure substances may be elements or compounds
pure substance - a substance that is made up of a single type of element or compound
distinguish between random and systematic errors; understand that the experimental design and procedure usually leads to systematic errors in measurement, which causes a deviation in a direction and appreciate that repeated trials and measurements will reduce random error but not systematic error
random = reliability systematic = validity
recall the relative energies of orbitals in energy levels to construct electron configurations for atoms and ions up to Z = 36 and recognise how the periodic table is arranged in relation to this [chapter 2]
s, p, d and f the periodic table is arranged into four blocks associated with the four sublevels Energy levels are divided into sublevels that have one or more orbitals. There are four types of orbitals, which have relative energies: s, p, d and f. Each orbital contains up to two electrons. The periodic table is based on the electron configurations of atoms and is arranged into four blocks associated with the four sublevels s, p, d and f. Na Z=11 → 1s^2, 2s^2, 2p^6, 3s^1 Cl Z=17 → 1s^2, 2s^2, 2p^6, 3s^2, 3p^5 Ar Z=18 → 1s^2, 2s^2, 2p^6, 3s^2, 3p^6
use the Arrhenius model to explain the behaviour of strong and weak acids and bases in aqueous solutions.
strong acids completely dissociate into hydronium ions in a solution Arrhenius acids increase the concentration of hydrogen ions in water. Arrhenius bases decrease the concentration of hydroxide ions in water.
construct and use appropriate graphical representations to communicate understanding, solve problems and make predictions; interpret graphs in terms of the relationship between the dependent and independent variables; draw and interpret best-fit lines or curves through data points, including evaluating when it can and cannot be considered as a linear function
yep
construct and use appropriate representations (including chemical symbols and formulas, and chemical and thermochemical equations) to communicate conceptual understanding, solve problems and make predictions
yep
construct and use appropriate representations, including ionic formulas, chemical formulas and chemical equations, to symbolise the reactions of acids and bases; and ionic equations to represent the reacting species and products in these reactions
yep
construct and use appropriate representations, including ionic formulas, chemical formulas, chemical equations and phase descriptions for chemical species to communicate conceptual understanding, solve problems and make predictions
yep
deduce and construct balanced chemical equations when reactants and products are specified and apply state symbols (s), (l), (g) and (aq)
yep
describe and explain how variations in the strength of the interactions between atoms, molecules or ions in the mobile and stationary phases can be used to separate components
yep
distinguish between precision and accuracy and appreciate that all measurements have limits to their precision and accuracy that must be considered when evaluating experimental results
yep
distinguish between qualitative and quantitative data; appreciate that quantitative data obtained from measurements is always associated with random error/measurement uncertainties
yep
explain the relationship between observable properties, including vapour pressure, melting point, boiling point and solubility, and the nature and strength of intermolecular forces, including dispersion forces, dipole-dipole attractions and hydrogen bonding within molecular covalent substances
yep
explain the relationship between the solubility of substances in water, including ionic and molecular substances, and the intermolecular forces between species in the substances and water molecules
yep
interpret, analyse and evaluate data and solubility curves to communicate conceptual understanding, solve problems and make predictions
yep
recall that chemical reactions and phase changes involve energy changes commonly observable as changes in the temperature of the surroundings and/or the emission of light
yep
recognise that changes in temperature can affect solubility and recall that most gases become less soluble as solvent temperature increases while most solutes become more soluble as the solvent temperature increases
yep
recognise that nano-materials are substances that contain particles in the size range 1-100 nm and have specific properties relating to the size of these particles
yep
understand that propagation of random error in data processing shows the impact of measurement uncertainties on the final result
yep
understand that the empirical formula expresses the simplest whole number ratio of elements in a compound
yep
understand that the number of significant figures in the result is based on the figures given in the data and determine the results of calculations to the appropriate number of significant figures
yep
use appropriate mathematical representation to solve problems and make predictions, including using the mole concept to calculate the mass of reactants and products; amount of substance in moles; number of representative particles; and molar mass of atoms, ions, molecules and formula units
yep
use appropriate mathematical representations to calculate the rate of chemical reactions by measuring the rate of formation of products or the depletion of reactants
yep
use molecular shape, understanding of symmetry, and comparison of the electronegativity of elements to explain and predict the polarity of molecules
yep
explain the relationship between the solubility of substances in water, including ionic and molecular substances, and the intermolecular forces between species in the substances and water molecules
yep - h-bonding
recognise that activation energy (Ea) is the minimum energy required for a chemical reaction to occur and is related to the strength and number of the existing chemical bonds; the magnitude of the activation energy influences the rate of a chemical reaction
yep - the minimum energy required to break the bonds - the higher the magnitude of the Ea the slower the rate of reaction
apply solubility rules to determine the products of reactions and to predict if a precipitate will form
yes
appreciate that the kinetic theory of gases applies to ideal gases and solve problems related to the ideal gas equation
yes
determine the presence of specific ions in solutions based on evidence derived from chemical reactions, including precipitation and acid-carbonate reactions
yes - precipitation is when salt is formed acid - base = salt and water acid - metal = salt and hydrogen acid - carbonate = salt + water + carbon dioxide
