26 KiB
HL Chemistry - A
The course code for this page is SCH3UZ.
1.1 - Review and base knowledge
Binary ionic and molecular compound nomenclature
An ionic bond is a bond formed between a metal cation and a non-metal anion. The compound name is written as the cation followed by the anion with the “-ide” ending. If the cation is multivalent (all transition metals aside from \(\text{Si}\) and \(\text{Co}\)), the charge of the cation should be written in parentheses in roman numerals between the two ions. \[\text{cation anion-ide || cation (roman charge) anion-ide}\]
!!! example - \(\text{NaCl}\): sodium chloride - \(\text{Al}_2\text{O}_3\): aluminium oxide - \(\text{CuO}\): copper (II) oxide - \(\text{Cu}_2\text{O}\): copper (I) oxide
A covalent bond is a bond formed between two non-metals. The compound name is written as the non-metal that appears first followed by the non-metal that appears second with the -ide ending, prefixing both with Greek prefixes to represent the number of atoms of that element. If there is only one atom of the first element, or if it would be unambiguous otherwise, “mono” is excluded for that element. The final vowel of the prefix is excluded if it ends with a vowel and the element following it starts with a vowel, except for the letter “i”. \[\text{prefix–element prefix–element–ide}\]
!!! example - \(\text{N}_2\text{O}_5\): dinitrogen pentoxide - \(\text{CO}\): carbon monoxide - \(\text{PI}_3\): phosphorous triiodide.
!!! warning The following common molecular compounds do not follow the rules above and have the following names instead:
- $\text{H}_2\text{O}_2$: hydrogen peroxide
- $\text{O}_3$: ozone
- $\text{H}_2\text{O}_\text{(l)}$: water
- $\text{CH}_4$: methane
- $\text{NH}_3$: ammoniaPolyatomic ions
Acids
All acids must be aqueous.
A binary acid consists of hydrogen atoms bonded to exactly one other non-metal element dissolved in water. It is written as the non-metal element prefixed by “hydro-” and suffixed with “-ic acid”. If the compound begins with a hydrogen atom but is not aqueous, it is a binary covalent molecule. \[\text{hydro–element–ic acid}\]
!!! example - \(\text{HCl}_\text{(aq)}\): hydrochloric acid - \(\text{HCl}\): hydrogen chloride - \(\text{H}_2{S}_\text{(aq)}\): hydrosulfuric acid
An oxyacid consists of a polyatomic ion bonded to hydrogen ions equal to its charge. It is written as the oxyacid with “-ate” replaced with “-ic”, “-ite” replaced with “-ous”, and with the suffix “acid”. \[\text{prefix-element-suffix acid}\]
!!! example - \(\text{H}_2\text{CO}_3\): carbonic acid - \(\text{HIO}_4\): periodic acid - \(\text{HIO}_3\): iodic acid - \(\text{HClO}_2\): chlorous acid - \(\text{HClO}\): hypochlorous acid
Hydrated salts
Types of reactions
Total and net ionic equations
2 - Atomic structure
!!! definition - The effective nuclear charge (\(Z_\text{eff}\)) is the net positive charge (attraction to the nucleus) experienced by an electron in an atom. - Electron shielding is decrease in the effective nuclear charge of an electron because of the repulsion of other electrons in lower-energy shells.
Atomic notation is used to represent individual atoms or ions. It is written in the form \(^M_Z \text{Symbol}^\text{Charge}\), where \(M\) is the mass number of the particle and \(Z\) is the atomic number of the particle.
!!! example - \(^1_1 \text{H}^{+}\) is the atomic notation for the most common hydrogen ion. - \(^{16}_8 \text{O}^{2-}\) is the atomic notation for the most common oxygen ion. - \(^{20}_{10} \text{Ne}\) is the atomic notation for the most common neon atom.
Isotopes
Isotopes are atoms of the same element but with different masses, or alternatively, atoms with the same number of protons but with different numbers of neutrons.
Radioisotopes are isotopes that are unstable (will spontaneously decay, are radioactive). Unstable atoms decay (break down) into one or more different isotopes of a different element. The half-life of a radioisotope is the time it takes for 50% of a sample’s atoms to decay.
!!! warning Radioisotopes are dangerous! They emit radiation, which is not at all good for human health in the vast majority of cases. However, there are also useful applications for radioisotopes today. For example, Cobalt-60 is used in radiation therapy to kill cancer tumours by damaging their DNA.
Atomic mass
The mass of every atom is represented relative to 1/12th of a carbon-12 atom. This mass is either unitless or expressed in terms of atomic mass units (amu or u). On the periodic table, the relative atomic mass (\(A_r\)) is shown, which is the sum of the masses of each isotope combined with their natural abundance (%abundance).
\[A_r = \text{%abundance}×\text{mass number of isotope}\] \[m_a = \Sigma A_r\]
When calculating the atomic mass from the graph from a mass spectrometer, the sum of the natural abundances of each isotope may not equal 100 or 1 (not in %abundance). In this case, calculation of %abundance will need to be done before solving for \(m_a\).
A mass spectrometer may also provide mass in the form of \(M/Z\), which represents mass over charge. For the sake of simplicity, \(Z=1\), so \(M/Z\) represents the mass of a particle.
Atomic radius
The atomic radius of an atom is the distance from the centre of the nucleus to approximately the outer boundary of the electron shell. This cannot be directly measured by scientists.
Ionisation energy
The first ionisation energy of an element is the minimum amount of energy required to remove one mole of electrons from one mole of gaseous atoms to form a mole of gaseous ions, so that \[\text{Q}_\text{(g)} \rightarrow \text{Q}_\text{(g)}^+ + \text{e}^-\]
Any subsequent ionisation energies of an element are the minimum amount of energy required to remove one additional mole of electrons. For example, the second ionisation energy would involve \[\text{Q}_\text{(g)}^+ \rightarrow \text{Q}_\text{(g)}^{2+} + \text{e}^-\]
It requires vastly more energy to remove an electron from a filled valence shell compared to an unfilled valence shell.
Electron affinity
The electron affinity of an atom is the amount of energy required or released to add an electron to a neutral gaseous atom to form a negative ion, such that \[\text{Q}_\text{(g)} + \text{e}^- \rightarrow \text{Q}^-_\text{(g)}\]
If energy is released, the atom has a negative electron affinity, and will form a stable ion.
If energy is required, the atom has a positive electron affinity, and will form an unstable ion (the ion will spontaneously decay).
Electronegativity
The electronegativity of an atom represents the ability of that atom to attract a pair of electrons in a covalent bond. It ranges from \(0\) to \(4\) on the Pauling scale. As electronegativity increases, the atom more strongly holds on to the electrons in its covalent bond, so the pair of electrons in that bond spend more time around the atom with the higher electronegativity.
Reactivity
The reactivity of an element is how “willing” it is to give up or gain electrons to fill its valence shell.
The reaction of an alkali metal with water always forms a hydroxide and hydrogen gas. For example, lithium reacts with water such that: \[2\text{Li}_\text{(s)} + 2\text{H}_2\text{O}_\text{(l)} \rightarrow 2\text{LiOH}_\text{(aq)} + \text{H}_{2 (g)}\]
The reaction of a halogen with hydrogen gas always forms a hydride. For example, fluorine reacts with hydrogen gas such that: \[\text{Fl}_\text{(g)} + \text{H}_\text{2 (g)} \rightarrow 2\text{HFl}_\text{(g)}\]
3 - Periodicity
Models
Please see SL Physics#Models for more information.
Periodic trends
Some trends in the periodic table include:
- atomic radius decreases when going across a period and increases when going down a group
- ionic radius decreases when going across a period for groups 1–13, then sharply increases and then increases for groups 15–17; it increases when going down a group
- electron affinity increases when going across a period and decreases when going down a group
- ionisation energy increases when going across a period and decreases when going down a group
- electronegativity increases when going across a period and decreases when going down a group
- reactivity of alkali metals increases when going down the group
- reactivity of halogens decreases when going down the group
When explaining these trends in the periodic table, it is best to use the following basic concepts to build on to larger points.
Across a period, the number of shells occupied by the electrons is the same but the number of protons in the nucleus increases. Therefore,
- the attraction of each electron to the nucleus (effective nuclear charge) increases as the number of protons increases
- shielding is unchanged as the number of electrons between the valence electrons and the nucleus is the same
Down a period, the number of shells occupied by the electrons increases, so valence electrons are further from the nucleus. Therefore,
- the attraction of valence electrons to the nucleus decreases due to the increasing distance
- shielding increases due to the increasing number of electrons between the valence electrons and the nucleus
!!! example To explain why there is a trend of decreasing atomic radius across a period:
- As the number of protons and electrons increase together, but the number of electron shells does not change, the effect of shielding remains unchanged while effective nuclear charge increases.
- This increased attraction to the nucleus reduces the atomic radius compared to other atoms before it.Period 3 oxides
Metal oxides (oxides of \(\text{Na}\) to \(\text{Al}\)) all form “giant ionic lattices” (alternatively just “lattices”) as an ionic bond is formed between a metal and a non-metal. These are typically solids because of their strong electrostatic attraction.
Metal oxides, except silicon (oxides of \(\text{P}\) to \(\text{Cl}\)) all form molecular compounds that exist as individual molecules as a covalent bond is formed between two non-metals. These are typically liquids or gases because of their weak intermolecular forces.
Period 3 oxides start basic but become more acidic when going across the period, with aluminium oxide being the turning point as an amphoteric substance (can be both an acid or a base). Basic oxides dissolve in water to form hydroxides while acidic oxides dissolve in water to form their respective oxyacids.
!!! example The following equations should be known by heart and are examples of period 3 oxides reacting with water. \[ \text{Na}_2\text{O}_\text{(s)} + \text{H}_2\text{O}_\text{(l)} → 2\text{NaOH}_\text{(aq)} \\ \text{MgO}_\text{(s)} + \text{H}_2\text{O}_\text{(l)} → \text{Mg(OH)}_\text{2 (aq)} \\ \text{P}_4\text{O}_\text{10 (s)} + 6\text{H}_2\text{O}_\text{(l)} → 4\text{H}_3\text{PO}_\text{4 (aq)} \\ \text{SO}_\text{3 (g)} + \text{H}_2\text{O}_\text{(l)} → \text{H}_2\text{SO}_\text{4 (aq)} \\ 3\text{NO}_\text{2 (g)} + \text{H}_2\text{O}_\text{(l)} → 2\text{HNO}_\text{3 (aq)} + \text{NO}_\text{(g)} \]
Alkali metals and halogens
The alkali metals are a family of highly reactive metals in group 1. They are very soft, and their melting and boiling points are relatively low, decreasing more when going down the group due to their weaker attraction. When reacted with water, they form hydrogen gas and metal hydroxides that have a high pH, hence the name “alkali” metals.
The halogens are a family of highly reactive non-metals in group 17. They occur diatomically (in molecules composed of two of the same element) and start as gases but become solids when going down the group due to stronger intermolecular forces. A single displacement reaction involving halogens only occurs if the more reactive halogen is not already bonded to the cation. Halogens are also very strong oxidising agents with their effectiveness increasing going up the group.
4.0 - Chemical bonding and structure
A chemical bond consists of the strong electronic interactions of the valence electrons between atoms that hold the atoms closer together. This only occurs if the atoms would reduce their potential energy by bonding.
!!! reminder When drawing a Lewis dot diagram, covalent bonds must be represented as two adjacent dots. When drawing a Lewis structure, covalent bonds must be represented as lines connecting the atoms.
If the process stage is required:
- Electrons destined to be shared must be encircled.
- Electrons to be transferred must have arrows pointing to their destination.
- x'es are used to represent additional electrons that have an unknown source.Nomenclature
Types of reactions
Total and net ionic equations
Percentage ionic character
Bonding is a spectrum. The percentage ionic character of a chemical bond shows roughly the amount of time valence electrons spend near an atom or ion in a bond. The difference between two elements’ electronegativity (ΔEN) indicates how covalent and how ionic the bond behaves.
If ΔEN is:
- less than 0.5, it behaves like a pure covalent bond
- between 0.5 and 1.7, it behaves like a polar covalent bond
- greater than 1.7, it behaves like an ionic bond
4.1 - Ionic bonding and structure
An ionic bond is the electrostatic attraction between oppositely charged ions. Electrons are transferred first, and then the bond forms via the attraction of the now-positive and negative ions. This reduces the potential energy of the ions and therefore increases their stability.
!!! definition Electrostatic attraction is the force of attraction between opposite charges.
!!! warning When expressing ionic bonds in a Lewis dot diagram, ions with charges of the same sign must never be placed next to one another.
Structure of ionic compounds
Ionic compounds are composed of a lattice structure (crystalline structure) of ions of alternating charges. A formula unit is the lowest ratio of positive to negative ions.
(Source:
Kognity)
!!! example In sodium chloride, the ratio of positive sodium ions to negative chloride ions is always 1:1, so its formula is NaCl.
In an ionic compound, the number of ions that each ion can touch is referred to as the coordination number. It is stated as “(cation)(anion) is (coordination number of cation):(coordination number of anion) coordinated”.
!!! example In the diagram above, each sodium ion touches six chloride ions, and each chloride ion touches six sodium ions. Therefore, “NaCl is 6:6 coordinated”.
4.2 - Covalent bonding
A covalent bond is the electrostatic attraction between pairs of valence electrons and nuclei. This causes atoms to “share” electrons instead of gaining or losing them. Covalent bonds form molecules, which in turn form molecular compounds (not covalent compounds).
(Source:
Kognity)
Whether a covalent bond is pure or polar indicates how evenly the shared electrons are shared between the atoms.
- A pure covalent bond has both nuclei attracting the valence electrons fairly evenly, so the difference in electronegativity (ΔEN) is low.
- A polar covalent bond has both nuclei attracting the valence electrons unevenly, so the ΔEN is high.
Bonding capacity
The bonding capacity of a non-metal describes the number of covalent bonds it can form.
Dative covalent bonds
Sometimes, one atom in a covalent bond may contribute both electrons in a shared pair.
4.3 - Covalent structures
Formal charge
There may be several correct ways to draw covalent bonds in Lewis structures and dot diagrams. Solving for the formal charge of each atom involved in a covalent bond can help identify the best structure to construct. The formal charge of an atom in a covalent bond represents the charge that that atom has. The sum of all formal charges in a covalently bonded compound is equal to the charge of the overall compound. The element with the lowest electronegativity is almost always in the centre.
The formal charge of an atom can be calculated using the following equation: \[\text{Formal charge} = \text{# of valence electrons of element} - \text{# of unpaired electrons} - \text{# of covalent bonds}\]
To find the best structure for a covalently bonded compound, the absolute value of the formal charge of all atoms in that compound should be minimised. Positively charged atoms will even accept dative covalent bonds from other atoms with negative formal charges.
!!! warning Some elements want formal charges of zero so much that they break the octet rule. These elements are \(\text{P, S, Cl, Br, I, and Xe}\).
Resonance structures
Even when considering formal charges, there may still be multiple best Lewis structures when molecules or polyatomic ions contain double or triple bonds. These equivalent structures are known as resonance structures, and the number of possible resonance structures is equal to the number of different positions for the double/triple bond. Double-sided arrows are used to show that the forms are resonant.
(Source:
Kognity)
The resonance structures of a compound show that none of the models is truly correct but instead the actual structure is somewhere in between all of them, and is not “flipping” between the various resonance structures.
!!! warning Molecules such as \(\text{SO}_2\) have resonance structures as the possible naive structures prior to involving formal charges are not considered to be resonant.
The resonance-hybrid structure shows that the actual strength of all three bonds is equal and somewhere between a single and double bond.
(Source:
Wikipedia)
Exceptions to the octet rule
Atoms such as boron and beryllium (\(\text{B}\) and \(\text{Be}\)) may form incomplete octets (less than 8 electrons) in their valence shell due to their status as small metalloids that form covalent bonds. In total, boron can sometimes need only 6 electrons while beryllium may have only 4 in their valence shells.
(Source:
Kognity)
!!! example \(\text{BeCl}_2\) and \(\text{BCl}_3\) exist.
Some elements in period 3 and beyond follow the formal charge exception above and may form expanded octets (more than 8 electrons and up to 12) in their valence shell. These include the aforementioned \(\text{P, S, Cl, Br, I, Xe}\), as well as metalloids.
Free radicals are molecules that end up with an odd number of electrons in their valence shell and are very reactive. Because one electron can never pair up with another, it remains forever alone.
??? example \(\text{NO}_2\) is a free radical as one of nitrogen’s atoms cannot pair with anything even after the formation of a dative covalent bond from oxygen.
Factors affecting bond strength
The strength of a bond is determined by the amount of energy required to break that bond (bond energy).
The length of a bond (bond length) has an inverse relationship with the strength of that bond, as the attraction of electrons to nuclei decreases with distance.
Multiple (double/triple) bonds are shorter than single bonds (a higher bond order) and are therefore stronger.
4.4 - Intermolecular forces
4.5 - Metallic bonding
11.1 - Uncertainties and errors in measurement and results
Please see SL Physics#Uncertainties and errors for more information.
11.2 - Graphical techniques
When plotting a graph:
- plot the independent variable on the horizontal axis and the dependent variable on the vertical axis
- label the axes, ensuring that the labels include units
- choose an appropriate scale for each axis
- give the graph an appropriate title at the bottom in title case
- draw a line of best fit
- include all uncertainties in the form of error bars
- if the error bars are too small to see, it should be noted explanation below
Titles
The title of a graph should clearly indicate what the graph represents under what conditions in title case, so that any onlooker should be able to identify the experiment. It should not include “vs.” Any legends present should be located under the graph.
??? example “Effect of Cat Deaths on Suicides in New Zealand”
Error bars
Please see SL Physics#Error bars for more information.
Line of best fit
Please see SL Physics#Uncertainty of gradient and intercepts for more information.
11.3 - Spectroscopic identification of organic compounds
Designing a scientific investigation
Scope
The scope of an experiment goes at the very beginning of it. It includes a general introduction to the topic of investigation as well as personal interest.
Research question
The research question of an experiment is a hyper-focused and specific question related to the topic. It contains and asks about the effect of an independent variable on a dependent variable.
Background information and hypothesis
!!! note This section can instead be placed immediately before the research question depending on the experiment.
In this section, scientific theories are provided to help the reader understand the rationale of the question, the design of the experiment, and data processing measures. If any theoretical/literature values are used, they should be introduced here.
A hypothesis consists of a justified prediction of the expected outcome and should be integrated with any background information.
Variables
!!! definition - The independent variable is the variable that is explicitly changed to attempt to affect the dependent variable. - The dependent variable is the variable that is directly monitored and measured in the experiment and is expected to change if the independent variable changes. - Controlled variables (also known as “control variables”) are variables that should be kept constant so they do not affect the dependent variable.
The independent variable, dependent variable, and any controlled variables should be listed under this section.
Materials
A list of materials and equipment should be listed here, as well as their precision. If a controlled variable needs to be measured, any instruments that would be used to do so should also be listed here.
Procedure
A clear, detailed, and concise set of instructions written in past tense should be placed in this section as either a numbered list or descriptive paragraph. To reduce confusion, if a numbered list is used, referring directly to numbers should be avoided, and referring to numbers recursively must never happen. A procedure must include:
- a clear, titled, labelled, and annotated diagram
- instructions for recording data (including for controlled variables)
If necessary, a “setup” section can be added as preparatory steps should not be listed in the main procedure.
Data collection
Data should be collected in an organised and titled table that should be prepared before the experiment. The data table must include:
- units with uncertainty, typically in the table header
- qualitative data (quantitative data can be optional in some experiments)
- repeated data/controlled variables, typically in the title
- any relevant information should be listed under the title
Only raw data prior to any processing or calculations, with the exception of averages, should be present in the data table.
A data table should be as concise as possible, and redundancy should be minimised. In that vein, trial numbers should not be recorded unless that data is relevant.
!!! example Table 1: Effect of Fat Content on Sugar Content in Ice Cream
| Fat Content (g ± 0.1 g) | Sugar Content (g ± 0.1 g) | Notes |
| --- | --- | --- |
| 2.0 | 5.1 | - strawberry ice cream |
| 0.1 | 2.3 | - mint chocolate chip ice cream |Whenever possible, data tables should not span multiple pages. If that is unavoidable, a new title with “…continued” and new column headers must be present at the top of each new page.
Data processing
A single sample calculation showing all steps should be present and clearly explained. The rest of the data can be processed without describing any steps. A single graph may be included if needed.
Some general rules include:
- units and uncertainties must be present in all calculations
- simple operations such as averages and conversions (e.g., g to kg) do not need to be explained
- the graph, if any, should span at least half of the page (ideally the full page) and should directly answer the research question
A final, reorganised, and processed data table should be present here, showing only relevant information.
Conclusion and evaluation
This section should be free of any new background information or calculations. It should, in sequence:
- summarise the results of the experiment without connecting it to the hypothesis
- identify whether the results of the experiment agree or disagree with the hypothesis
- evaluate 3–5 systematic errors (usually) present in the experiment, both in the procedure and in data collection/processing, in decreasing order of impact to the experiment
The evaluation of systematic errors should include:
- a description of the error
- how the error affected the data
- how the error affected the final result
- how the error can be remedied with available school resources