Atomic Structure and the Periodic Tabl
2.2.1 Describe the structure of the atom as a central nucleus containing neutrons and protons surrounded by electrons in shells.
Atoms consist of a dense, positively charged nucleus at the center.
The nucleus contains two types of subatomic particles:
Protons: Positively charged particles.
Neutrons: Electrically neutral particles.
Negatively charged particles called electrons orbit the nucleus in specific energy levels or shells.
The first electron shell can hold up to 2 electrons.
The second and third electron shells can each hold up to 8 electrons.
2.2.2 State the relative charges and relative masses of a proton, a neutron, and an electron.
| Particle | Relative Mass | Relative Charge |
Proton. | 1 mass unit | +1 | |
Neutron | 1 mass unit | 0 |
Electron | negligible | -1 |
2.2.3 Define proton number / atomic number (the smaller number represented on the periodic table) as the number of protons in the nucleus of an atom.
The proton number (Z), also known as the atomic number, is a fundamental characteristic of an element.
It represents the number of protons present in the nucleus of each atom of that element.
The atomic number determines the element's identity and its position on the periodic table.
2.2.4 Define mass number / nucleon number (the larger number represented on the periodic table) as the total number of protons and neutrons in the nucleus of an atom.
The mass number (A), also known as the nucleon number, is the sum of the number of protons and neutrons in the nucleus of an atom.
Number of neutrons = Mass number (A) - Atomic number (Z).
2.2.5 Determine the electronic configuration of elements and their ions with proton number 1 to 20, e.g. 2,8,3.
The electronic configuration describes the arrangement of electrons in the different energy levels or shells around the nucleus of an atom or ion.
Electrons fill the shells starting from the innermost shell.
For elements with a proton number of 1 to 20:
The first shell (n=1) can hold a maximum of 2 electrons.
The second shell (n=2) can hold a maximum of 8 electrons.
The third shell (n=3) starts filling after the first two are full and can hold up to 8 electrons for these elements.
Ions are formed when atoms gain or lose electrons to achieve a stable electron configuration. Positive ions (cations) have lost electrons, and negative ions (anions) have gained electrons.
II. Isotopes:
2.2.6 State that:
(a) Group VIII noble gases have a full outer shell.
Noble gases (Group 18 or VIII) have a stable electron configuration with a complete outer shell of electrons (usually 8, except for Helium which has 2). This full outer shell makes them chemically unreactive.
(b) the number of outer shell electrons is equal to the group number in Groups I to VII.
For elements in Groups 1 to 17 (I to VII), the number of electrons in the outermost shell (valence electrons) is equal to the group number. For example, Group 1 elements have 1 valence electron, Group 17 elements have 7 valence electrons.
(c) the period number of an element indicates the number of electron shells occupied by electrons.
The horizontal rows in the periodic table are called periods. The period number corresponds to the number of electron shells that contain electrons in an atom of that element. For example, elements in Period 3 have electrons in three energy levels.
2.3.1 Define isotopes as atoms of the same element that have the same number of protons but different numbers of neutrons.
Isotopes are atoms of the same element that have the same atomic number (number of protons) but different mass numbers (due to a different number of neutrons).
For example, Carbon-12 (12C), Carbon-13 (13C), and Carbon-14 (14C) are isotopes of carbon. They all have 6 protons but have 6, 7, and 8 neutrons, respectively.
2.3.2 Interpret and use symbols for atoms, e.g. 612C and ions, e.g. 1735Cl−
The notation for an atom or ion is generally represented as: ZAXq where:
X is the chemical symbol of the element.
A is the mass number (number of protons + neutrons).
Z is the atomic number (number of protons).
q is the charge of the ion (if present).
For example:
612C: Carbon atom with a mass number of 12 and an atomic number of 6 (6 protons, 6 neutrons, 6 electrons).
1735Cl−: Chloride ion with a mass number of 35 and an atomic number of 17 (17 protons, 18 neutrons, 18 electrons due to the -1 charge).
2.3.3 State that isotopes of the same element have the same chemical properties because they have the same number of electrons and therefore the same electronic configuration.
The chemical properties of an element are primarily determined by the number and arrangement of its electrons, particularly the valence electrons.
Since isotopes of the same element have the same number of protons and therefore the same number of electrons (in a neutral atom), they have the same electronic configuration.
Consequently, isotopes of the same element exhibit essentially the same chemical behavior and undergo the same types of chemical reactions.
2.3.4 Calculate the relative atomic mass of an element from the relative masses and abundances of its isotopes.
The relative atomic mass (Ar) of an element is the weighted average of the relative masses of its naturally occurring isotopes, taking into account their relative abundances.
The formula for calculating relative atomic mass is: Ar=total abundance(mass of isotope 1×abundance of isotope 1)+(mass of isotope 2×abundance of isotope 2)+...
Abundance is usually given as a percentage, in which case the total abundance is 100.
III. The Periodic Table: Group Properties and Trends:
8.2.1 Describe the Group I alkali metals, lithium, sodium and potassium, as relatively soft metals with general trends down the group, limited to: (a) decreasing melting point (b) increasing density (c) increasing reactivity
Group I Alkali Metals (Lithium, Sodium, Potassium):
Are relatively soft metals that can be easily cut with a knife.
Exhibit the following general trends as you go down the group:
(a) Decreasing Melting Point: The forces of attraction between atoms weaken, requiring less energy to break.
(b) Increasing Density: The increase in mass is more significant than the increase in atomic size.
(c) Increasing Reactivity: The outermost electron is further from the nucleus and more easily lost, leading to a greater tendency to form positive ions.
8.2.2 Predict the properties of other elements in Group I, given information about the elements.
Based on the trends observed in Lithium, Sodium, and Potassium, it is possible to predict the properties of other alkali metals (e.g., Rubidium, Cesium, Francium). For instance, we can expect them to be even softer, have lower melting points, higher densities (initially), and be more reactive than potassium.
8.3.1 Describe the Group VII halogens, chlorine, bromine and iodine, as diatomic non-metals with general trends down the group, limited to: (a) increasing density (b) decreasing reactivity
Group VII Halogens (Chlorine, Bromine, Iodine):
Exist as diatomic molecules (e.g., Cl2, Br2, I2).
Are non-metals.
Exhibit the following general trends as you go down the group:
(a) Increasing Density: The increase in mass is more significant than the increase in atomic size.
(b) Decreasing Reactivity: The outermost shell becomes further from the nucleus, making it harder to attract an electron to complete the outer shell.
8.3.2 State the appearance of the halogens at room temperature and pressure (r.t.p.) as: (a) chlorine, a pale yellow-green gas (b) bromine, a red-brown liquid (c) iodine, a grey-black solid
At room temperature and pressure:
(a) Chlorine (Cl2) is a pale yellow-green gas.
(b) Bromine (Br2) is a red-brown liquid.
(c) Iodine (I2) is a grey-black solid.
8.3.3 Describe and explain the displacement reactions of halogens with other halide ions, including what you would observe, and the word equations for each reaction.
A more reactive halogen can displace a less reactive halide ion from its aqueous solution.
Reactivity order of halogens: Chlorine > Bromine > Iodine.
Examples:
Chlorine water + Potassium bromide solution → Potassium chloride solution + Bromine
Observation: The colorless potassium bromide solution turns orange-brown as bromine is formed.
Word equation: Chlorine + Potassium bromide → Potassium chloride + Bromine
Chlorine water + Potassium iodide solution → Potassium chloride solution + Iodine
Observation: The colorless potassium iodide solution turns brown as iodine is formed.
Word equation: Chlorine + Potassium iodide → Potassium chloride + Iodine
Bromine water + Potassium iodide solution → Potassium bromide solution + Iodine
Observation: The colorless potassium iodide solution turns brown as iodine is formed.
Word equation: Bromine + Potassium iodide → Potassium bromide + Iodine
A less reactive halogen cannot displace a more reactive halide ion (e.g., Bromine cannot displace chloride from a potassium chloride solution).
8.3.4 Predict the properties of other elements in Group VII, given information about the elements.
Based on the trends in Chlorine, Bromine, and Iodine, we can predict properties of other halogens (e.g., Fluorine, Astatine). Fluorine is expected to be a pale yellow gas and the most reactive halogen, while Astatine is predicted to be a dark solid and the least reactive.
8.4.1 Describe the transition elements as metals that: (a) have high densities (b) form coloured compounds (c) often act as catalysts in chemical reactions.
Transition Elements:
Are located in the central block of the periodic table (Groups 3-12).
Exhibit typical metallic properties, including:
(a) High Densities: Generally have higher densities compared to Group I and II metals.
(b) Form Coloured Compounds: Their ions often form solutions and compounds that are brightly coloured due to the presence of partially filled d orbitals.
(c) Often Act as Catalysts: Many transition metals and their compounds can act as catalysts, speeding up chemical reactions without being used up themselves (e.g., iron in the Haber process, nickel in hydrogenation).
8.4.2 Describe transition elements as having ions with variable oxidation numbers, including iron(II) and iron(III).
Variable Oxidation Numbers: Transition elements can exhibit more than one oxidation state (or valency) because they can lose electrons from both their s and d subshells.
Iron Ions: Iron is a common transition metal that exists in two common ionic forms:
Iron(II) ions (Fe2+): Formed when iron atoms lose two electrons.
Iron(III) ions (Fe3+): Formed when iron atoms lose three electrons.
8.5.1 Describe the Group VIII noble gases as unreactive, monatomic gases and explain this in terms of their electronic configuration.
Group VIII Noble Gases (Helium, Neon, Argon, etc.):
Are chemically unreactive or inert.
Exist as monatomic gases (single atoms) because they have very weak interatomic forces.
Electronic Configuration: Their unreactivity is due to their stable electronic configuration with a full outer shell of electrons (2 for Helium, 8 for the rest). This stable arrangement means they have little or no tendency to gain, lose, or share electrons to form chemical bonds.