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2.1 Atoms and Reactions - Coggle Diagram

- - - - Protons and neutrons, collectively called nucleons, are located in the atom's nucleus. Protons have a positive charge, while neutrons are electrically neutral. The nucleus is extremely small and dense, containing the majority of an atom's mass.
        .
      - Electrons, which have a negative charge, orbit the nucleus in shells or orbitals. Although electrons occupy a large region of space around the nucleus, most of an atom's volume is actually empty space
      - Protons
        
        Relative Charge: +1
        
        Relative Mass: 1
      - Neutron
        
        Relative Charge: 0
        
        Relative Mass: 1
      - Electron
        
        Relative Charge: -1
        
        Relative Mass: 0.0005
    - - Nuclear symbols provide a concise way to represent the composition of an atom. They indicate the number of protons, neutrons, and electrons in an atom using two key numbers:
        
        Mass number (A) - The total number of nucleons (protons and neutrons) in the nucleus.
        .
        
        Atomic number (Z) - The number of protons in the nucleus, which uniquely identifies the element
      - To determine the number of protons, electrons, and neutrons in a neutral atom:
        
        The number of protons is equal to the atomic number (Z).
        
        In a neutral atom, the number of electrons is equal to the number of protons.
        
        The number of neutrons is calculated by subtracting the atomic number (Z) from the mass number (A).
    - - Ions are atoms that have gained or lost electrons, resulting in an overall positive or negative charge.
      - Negative ions (anions) have more electrons than protons. For example, a particle with 1 proton and 2 electrons would be a negative ion with an overall charge of -1.
        .
      - Positive ions (cations) have fewer electrons than protons. For example, a particle with 3 protons and 2 electrons would be a positive ion with an overall charge of +1
      - To determine the number of protons, electrons, and neutrons in an ion:
        
        The number of protons is equal to the atomic number.
        
        The number of electrons is calculated by subtracting the ion's charge from the atomic number (for cations) or adding the ion's charge to the atomic number (for anions).
        
        The number of neutrons is calculated by subtracting the atomic number from the mass number.
  - - - Our understanding of the atom's structure has grown considerably over the past few centuries. As scientists have discovered new experimental evidence, they've needed to revise their atomic models to more accurately explain their findings.
      - It's crucial to recognise that the currently accepted model is still just a model - a simplified depiction of reality meant to help us understand chemical phenomena.
    - - Dalton's solid sphere model (early 19th century)
        
        Proposed by John Dalton, this model suggested that atoms were indivisible, solid spheres.
        
        It also suggested that different elements were made up of different types of these spheres.
      - Thomson's plum pudding model (1897)
        
        The discovery of the electron by J.J. Thomson proved that atoms were divisible.
        
        His 'plum pudding' model depicted atoms as spheres of positive charge with negative electrons randomly embedded throughout, similar to the raisins in a plum pudding.
    - - The year 1909 saw Ernest Rutherford, along with his students Hans Geiger and Ernest Marsden, perform a ground-breaking experiment that would redefine atomic theory.
      - Involved firing alpha particles (which are positively charged) at a sheet of thin gold foil inside a vacuum.
      - According to the plum pudding model, the positive 'pudding' would slightly deflect most of the particles.
      - Contrary to expectations, most alpha particles passed straight through, with a few deflecting at large angles or even bouncing back.
      - Rutherford's conclusions:
        
        The fact that most alpha particles passed through the gold foil undeflected indicated that atoms are mostly empty space.
        
        The small number of particles that deflected at large angles or bounced back suggested the presence of a tiny, dense, positively charged nucleus at the centre of the atom. This nucleus was responsible for the strong deflection of the positively charged alpha particles due to electrostatic repulsion.
        
        Since the atom was mostly empty space with a dense nucleus, Rutherford proposed that negative electrons must orbit the nucleus in a 'cloud' to maintain the atom's overall neutrality.
      - This discovery, based on the surprising results of the gold foil experiment, gave birth to the nuclear model of the atom.
    - - Rutherford's nuclear model was a significant step forward but required adjustments as further evidence came to light:
        
        Discovery of the proton
        
        Research by Henry Moseley showed that the nuclear charge increases by one unit from one element to the next.
        
        This insight led Rutherford to identify positively charged particles within the nucleus, which he named protons.
        
        The differences in nuclear charges among elements could now be explained by the presence of varying numbers of protons.
        
        Prediction and discovery of the neutron
        
        A remaining mystery was why atomic nuclei were heavier than they should be if they contained only protons.
        
        Rutherford speculated about the existence of another type of nuclear particle.
        
        James Chadwick later discovered the neutron, which provided the mass needed to account for the discrepancy in the nucleus's weight.
    - - A flaw in Rutherford's model was its inability to explain why electrons, attracted to the nucleus, wouldn't simply collapse into it.
        
        To solve this, Niels Bohr suggested a model in which:
        
        Electrons orbit the nucleus in specific energy levels, or shells.
        
        They can jump between these shells by absorbing or emitting light of particular wavelengths.
        
        This model accounted for the observed spectra of light emitted or absorbed by atoms, a phenomenon Rutherford's model could not explain.
    - - Subsequent experiments revealed that electrons within the same shell could possess different energy levels, leading to the introduction of subshells in the Bohr model.
      - While even more accurate atomic models exist today, the refined Bohr model remains useful due to its simplicity and ability to explain many chemical concepts, such as trends in bonding and ionisation energy.
  - - - Elements are defined as:
        
        Substance made up of only one type of atom, where all the atoms of a substance have the same number of protons.
      - Isotopes are defined as:
        
        Atoms of the same element that have the same number of protons but different numbers of neutrons.
        
        Most elements are made up of a mixture of isotopes.
    - - Chemical properties:
        
        Isotopes have the same electron configuration.
        
        This means they have the same chemical properties (e.g. reactivity
      - Physical properties:
        
        Isotopes have slightly different physical properties (e.g. mass and density).
        
        This is because physical properties depend on atomic mass.
    - - Atoms have tiny absolute masses that are impractical to measure directly. Instead, the mass of one atom is compared to 1/12th the mass of a carbon-12 atom. This is the atom's relative mass.
      - Key definitions:
        
        Relative isotopic mass - The mass of an atom of an isotope compared to 1/12 the mass of a carbon-12 atom.
        
        Relative isotopic mass is specific to a particular isotope of an element.
        
        It is usually very close to being a whole number.
        
        At A level, relative isotopic masses are rounded to one decimal place.
        
        E.g. The relative isotopic masses of the two main chlorine isotopes (35Cl and 37Cl) are 35.0 and 37.0 respectively.
        
        Relative atomic mass (Ar) - The weighted mean mass of an atom of an element compared to 1/12 the mass of a carbon-12 atom.
        
        Ar is an average of the masses of all the isotopes of an element, weighted according to their natural abundances.
        
        Since it's an average, Ar is not usually a whole number.
        
        The next section explains how to calculate relative atomic mass from isotopic abundances.
        
        Relative molecular mass (Mr) - The average mass of a molecule compared to 1/12 the mass of a carbon-12 atom.
        
        For simple covalent molecules, Mr is calculated by adding up the Ar values of all the atoms in one molecule.
        
        Relative formula mass - Used for ionic compounds. It is calculated by adding up the Ar values of all the ions in one formula unit.
  - - - A mass spectrum plots the relative abundance of ions against their mass-to-charge ratio (m/z).
      - The x-axis displays the m/z values. The m/z of each peak equals the relative mass of the ion.
      - The y-axis indicates the relative abundance of each ion, either in arbitrary units or as a percentage.
      - For elemental samples, each peak represents a different isotope.
        
        The relative height of each peak shows how abundant the isotope is.
        
        Elements with only one stable isotope will show a single peak.
      - For molecular samples, the peak at the highest m/z value is the molecular ion (M+).
        
        The m/z of the molecular ion peak matches the molecule's relative molecular mass (Mr).
        
        Peaks at lower m/z values come from fragments of the molecular ion.
- - - - An acid is defined as a substance that donates protons (H+) when dissolved in water. These protons combine with water molecules to form hydronium ions (H3O+), indicating the acidic nature of the solution.
        
        Examples of Common Acids include:
        
        Hydrochloric acid (HCl)
        
        Sulfuric acid (H2SO4)
        
        Nitric acid (HNO3)
        
        Ethanoic acid (CH3COOH)
      - A base, in contrast, is a substance that accepts protons. When a base dissolves in water, it can produce hydroxide ions (OH-), making the solution alkaline. Bases that dissolve in water are specifically called alkalis.
        
        Examples of common alkalis include:
        
        Sodium hydroxide (NaOH)
        
        Potassium hydroxide (KOH)
        
        Aqueous ammonia (NH3 in H2O)
      - Aqueous ammonia is a special case; it doesn’t produce hydroxide ions directly but accepts protons from water to form ammonium (NH4+) and hydroxide ions (OH-):
        NH3 + H2O ➔ NH4+ + OH-
    - - The strength of an acid or a base is determined by its ability to dissociate in water. This dissociation process can be reversible, as shown in the equations below:
        
        Acid dissociation: HA + H2O ⇌ H3O+ + A-
        Base dissociation: B + H2O ⇌ BH+ + OH-
      - Strong acids and bases:
        
        Strong acids, such as hydrochloric acid (HCl), and strong bases, like sodium hydroxide (NaOH), dissociate completely in water. This results in a significant release of H+ and OH- ions, respectively, with the forward reaction being predominantly favoured:
        
        HCl(aq) ➔ H+(aq) + Cl-(aq)
        NaOH(aq) ➔ Na+(aq) + OH-(aq)
      - Weak acids and bases:
        
        Conversely, weak acids and bases, like ethanoic acid (CH3COOH) and ammonia (NH3), only partially dissociate in water, releasing fewer H+ and OH- ions. In these cases, the reverse reaction is favoured:
        
        CH3COOH(aq) ⇌ H+(aq) + CH3COO-(aq)
        NH3(aq) + H2O(l) ⇌ NH4+(aq) + OH-(aq)
    - - Neutralisation is the reaction between an acid and an alkali, resulting in the formation of water and a salt. This reaction specifically involves the combination of H+ ions from the acid with OH- ions from the alkali to form water (H2O).
        
        The general ionic equation for this process is:
        
        H+(aq) + OH-(aq) ➔ H2O(l)
      - A salt is an ionic compound formed when the H+ ions in an acid are replaced by metal ions or other positive ions, such as ammonium ions (NH4+). In a neutralisation reaction, the salt is comprised of the cation from the base, which can be a metal or ammonium ion, and the anion from the acid.
      - The type of salt produced depends on the specific acid used in the reaction:
        
        Sulfuric acid (H2SO4)
        
        Anion in Salt- Sulfate (SO42-)
        
        Hydrochloric acid (HCl)
        
        Anion in Salt- Chloride (Cl-)
        
        Nitric acid (HNO3)
        
        Anion in Salt- Nitrate (NO3-)
    - - Acids and reactive metals
        
        Acids react with certain metals, producing a salt and hydrogen gas:
        
        Metal + acid ➔ salt + hydrogen
        
        For example, calcium reacting with sulfuric acid: Ca(s) + H2SO4(aq) ➔ CaSO4(aq) + H2(g)
      - Acids and metal oxides
        
        These reactions yield a salt and water:
        
        Metal oxide + acid ➔ salt + water
        
        An example is zinc oxide with hydrochloric acid: ZnO(s) + 2HCl(aq) ➔ ZnCl2(aq) + H2O(l)
      - Acids and metal hydroxides
        
        Similar to oxides, these reactions produce a salt and water:
        
        Metal hydroxide + acid ➔ salt + water
        
        For example, potassium hydroxide reacting with nitric acid: KOH(aq) + HNO3(aq) ➔ KNO3(aq) + H2O(l)
      - Acids and metal carbonates
        
        These reactions produce a salt, water, and carbon dioxide:
        
        Metal carbonate + acid ➔ salt + water + carbon dioxide
        
        An example is sodium carbonate with hydrochloric acid: Na2CO3(s) + 2HCl(aq) ➔ 2NaCl(aq) + H2O(l) + CO2(g)
  - - - Titration is an analytical method used to determine the concentration of a solution of known volume by gradually adding a solution of known concentration to it until neutralisation occurs.
      - A solution with an unknown volume and concentration is referred to as the analyte. It is measured using a pipette and placed into a conical flask.
      - The standard solution, which has a known concentration, is filled into a burette and slowly added to the analyte.
      - To visually determine when neutralisation has occurred, an indicator is added to the flask, causing a colour change at the end point - the point at which the reaction has just reached completion and neither reactant is in excess.
    - - Measure the analyte into a conical flask using a pipette, and add an indicator.
      - Fill the burette with the standard solution and note the initial reading.
      - Gradually add the standard solution to the flask, swirling to mix, until you approximately reach the point of neutralisation.
      - Refill the burette and repeat steps 2 and 3 to get an approximate idea of where neutralisation occurs.
      - Conduct a precise titration by adding the standard solution slowly to identify the exact end point. Record the final burette reading, also known as the titre volume.
      - Perform the titration multiple times to obtain concordant results, which are titre volumes that differ by no more than 0.10 cm3.
      - Calculate the mean titre volume using the concordant results.
    - - When reading a 50 cm3 burette, the volume should be recorded to the nearest half division (±0.05 cm3), with the bottom of the meniscus either on a mark or between two marks.
      - The meniscus is the curved surface of the liquid in the burette, which is caused by the adhesive forces between the liquid and the glass.
      - To ensure accurate readings, always read the volume at the bottom of the meniscus, as shown in the image below.
    - - Indicators are crucial for visually identifying the end point of a titration through a colour change, indicating neutralisation.
        
        Common indicators include:
        
        Methyl orange, which changes from yellow to red when adding acid to alkali.
        
        Phenolphthalein, which changes from pink to colourless when adding acid to alkali.
        
        Conduct titrations against a white background to easily observe the colour change.
    - - To determine an unknown concentration via titration, it's essential to use a standard solution with a precisely known concentration.
        
        These solutions are prepared by:
        
        Weighing a precise mass of solute using a balance.
        
        Dissolving the weighed solute in distilled water in a beaker.
        
        Transferring the solution to a volumetric flask of a specific volume.
        
        Rinsing the beaker several times with distilled water and adding the rinsings to the volumetric flask.
        
        Diluting the solution to the mark on the volumetric flask with a pipette.
        
        Inverting the flask several times to ensure thorough mixing of the solution.
  - - - Oxidation numbers, also known as oxidation states, are assigned to each atom or ion in a compound to indicate the degree of oxidation. These numbers represent the total number of electrons transferred to or from an atom within the compound.
      - Rules for assigning oxidation numbers:
        
        Uncombined elements (e.g. Ne) have an oxidation number of 0.
        
        Atoms in molecules bonded to identical atoms (e.g. O2) also have an oxidation number of 0.
        
        A monatomic ion's oxidation number is equal to its charge. E.g., Fe3+ has an oxidation number of +3.
        
        In a molecular ion, the sum of the oxidation numbers equals the overall charge. E.g., in SO42-, S is +6 and O is -2, summing to -2.
        
        In a neutral compound, the sum of the oxidation numbers is 0. E.g., in Al2O3, Al is +3 and O is -2, summing to 0.
        
        Combined oxygen usually has an oxidation number of -2, except in peroxides, where it is -1, and in OF2 where it is +2.
        
        Combined hydrogen usually has an oxidation number of +1, except in metal hydrides, where it is -1.
      - In a compound or ion containing 2 non-metals, the negative oxidation number is given to the more electronegative element. Fluorine always has an oxidation number of -1 because it is the most electronegative element.
    - - When elements can exhibit multiple oxidation numbers, Roman numerals are used after the element's name to specify its oxidation number.
      - For example, iron(II) signifies iron with an oxidation number of +2, while iron(III) indicates an oxidation number of +3.
    - - Ions ending in '-ate' are made up of oxygen and another element. For example, sulfates include sulfur, while nitrates include nitrogen.
      - If the element bonded to oxygen has variable oxidation numbers, Roman numerals are used after the compound's name to differentiate between these numbers.
      - Sulfate(VI)
        
        Oxidation Number: +6
        
        Formula of ion: SO42-
      - Sulfate(IV)
        
        Oxidation Number: +4
        
        Formula of ion: SO32-
      - Nitrate(III)
        
        Oxidation Number: +3
        
        Formula of ion: NO2-
  - - - Oxidation is the process of electron loss. Reduction is the process of electron gain.
      - These processes occur simultaneously, which is why they are referred to collectively as a "redox reaction".
      - An oxidising agent accepts electrons and undergoes reduction itself.
      - A reducing agent donates electrons and undergoes oxidation itself.
      - For example, consider the reaction between sodium and chlorine to form sodium chloride:
        
        Sodium is oxidised because it loses electrons, acting as the reducing agent.
        
        Chlorine is reduced because it gains electrons, acting as the oxidising agent.
    - - Oxidation numbers increase by 1 for every electron lost.
      - Oxidation numbers decrease by 1 for every electron gained.
      - To determine if a redox reaction has occurred, compare the oxidation numbers before and after the reaction:
        
        An increase indicates oxidation
        
        A decrease indicates reduction
      - For example, in the reaction:
        
        Na(s) + 1⁄2Cl2(g) ➔ NaCl(s)
        
        Sodium is oxidised because its oxidation number increases from 0 (in Na) to +1 (in NaCl).
        
        Chlorine is reduced because its oxidation number decreases from 0 (in Cl2) to -1 (in NaCl).
    - - A reaction between a metal and an acid to produce a salt and hydrogen gas exemplifies a redox reaction:
        
        Metal atoms lose electrons and are oxidised.
        
        Hydrogen ions gain electrons and are reduced.
      - For example:
        
        Mg(s) + 2HCl(aq) ➔ MgCl2(aq) + H2(g)
        
        Magnesium is oxidised because its oxidation number increases from 0 to +2.
        
        Hydrogen ions are reduced because their oxidation number decreases from +1 to 0.
        
        Chloride ions are neither oxidised nor reduced because their oxidation number remains unchanged. Any species whose oxidation number remains constant throughout the reaction acts as a spectator ion and is not included in the ionic equation:
        
        Mg(s) + 2H+(aq) ➔ Mg2+(aq) + H2(g)