Fiche de révision : Introduction to Atomic and Periodic Chemistry

📋 Course Outline

1. Atomic structure and isotopes
2. Periodic table and electron configuration
3. Ionic compounds and formulae
4. Relative atomic mass and molecular mass
5. Chemical equations and balancing
6. Group 1 metals and Group 7 halogens
7. Ionic tests and flame colours
8. Noble gases and their uses

📖 1. Atomic structure and isotopes

🔑 Key Concepts & Definitions

• Atomic nucleus : The atomic nucleus is the central part of an atom containing protons and neutrons.
• Atomic number : The atomic number is the number of protons in an atom and fixes which element it is.
• Mass number : The mass number is the total number of protons plus neutrons in an atom.
• Isotopes : Isotopes are atoms of the same element with the same atomic number but different numbers of neutrons.

📝 Essential Points

• Protons and neutrons each have a mass of 1 amu, while electrons have a negligible mass about 2000 times smaller.
• Protons have charge +1 and neutrons have charge 0, so the nucleus has an overall positive charge.
• An atom has no overall charge because the number of electrons equals the number of protons.
• Ions form when an atom gains or loses electrons, giving a negative ion after gaining and a positive ion after losing.
• For any atom or simple ion: atomic number = number of protons, and mass number = protons + neutrons, letting you find electron count from the ion charge.
• Relative atomic mass for an element with more than one isotope is calculated using the isotopes’ mass numbers together with their relative abundances.

💡 Memory Hook

Nucleus = (p + n) with charge from p only; isotopes change n, not p (same element, different neutrons).

📖 2. Periodic table and electron configuration

🔑 Key Concepts & Definitions

• Group number : Group number is the position of an element in the periodic table and matches the number of electrons in its outer shell.
• Period number : Period number is the row position of an element and matches the number of occupied electron shells.
• Electron shell capacity : Electron shell capacity is the maximum number of electrons a shell can hold, given by 2 in the first shell, 8 in the second, and 8 then 18 in higher shells as listed.

📝 Essential Points

• Elements in the same group have similar chemical properties because they have the same number of outer-shell electrons.
• In the periodic table, the elements are arranged in increasing atomic number from left to right and top to bottom.
• Group 1 becomes more reactive when moving down the group, and Group 1 boiling and melting points decrease down the group (lithium does not melt when expected because it does not react with water).
• Group 7 becomes less reactive when moving down the group, and Group 7 boiling and melting points increase down the group.
• Many reactions involve losing or gaining electrons to form ions with charges that result from changing electron number.
• Group 0 elements are very unreactive because their outer shell is full, so they have little tendency to gain or lose electrons.

💡 Memory Hook

Group 1: down = more reactive + lower m.p./b.p.; Group 7: down = less reactive + higher m.p./b.p.

📖 3. Ionic compounds and formulae

🔑 Key Concepts & Definitions

• Positive ions : Positive ions are ions with a net positive charge formed by metals losing electrons.
• Negative ions : Negative ions are ions with a net negative charge formed when non-metals gain electrons or when non-metal groups form anions.
• Ionic compound formula : An ionic compound formula is written using the ion symbols with subscripts that give an overall neutral total charge.
• Roman numeral metal charge : A Roman numeral in a metal’s ion name indicates the metal’s positive charge used in the compound.

📝 Essential Points

• Ionic compounds must be neutral overall, so the total positive charge equals the total negative charge in the formula.
• When a metal can form more than one positive ion, write its charge as a Roman numeral beside the metal name.
• Negative ion names often end with -id, for example bromine becomes bromid(e), and fluoride becomes fflworid(e).
• If a non-metal ion includes oxygen, its negative-ion name can follow the pattern ending in -ad, as shown by nitrad from NO3- (nitrogen + oxygen).
• To write a formula: find ion symbols and charges, count positive and negative ions, then balance them so charges cancel completely.
• Example: sodium chloride is Na+ and Cl- with equal charges, so its formula is NaCl.

💡 Memory Hook

Neutral compounds balance charges: total + equals total −, so ion counts must make the charges cancel.

📖 4. Relative atomic mass and molecular mass

🔑 Key Concepts & Definitions

• Relative atomic mass Ar : Relative atomic mass is the weighted average mass of an atom, taking account of its isotopes and how common they are.
• Isotopic abundance percentage : Isotopic abundance is the relative percentage of each isotope in a naturally occurring sample used to find Ar.
• Relative formula mass Mr : Relative formula mass is the total relative mass of all atoms in one formula unit, found by adding the Ar values for each atom.

📝 Essential Points

• Ar is calculated from isotopic abundances using $($.
• For chlorine, using 75% of $^{35}\text{Cl}$ and 25% of $^{37}\text{Cl}$ gives $Ar=35.5$.
• $Mr$ for a molecular formula is the sum of the relevant atom Ar values, e.g. $Mr(\text{NaCl})=23+35.5=58.5$.
• $Mr$ for $\text{Na}_2\text{CO}_3$ uses the multipliers in the formula, giving $Mr=2\times 23+12+3\times 16=106$.
• For any isotope of an element, its Ar value equals its mass number from the periodic table.

💡 Memory Hook

Ar = weighted average: add (abundance × isotope mass) ÷ 100; then $Mr$ = sum of Ar’s in the formula.

📖 5. Chemical equations and balancing

🔑 Key Concepts & Definitions

• Chemical equation : A chemical equation is a way to write a reaction using words or symbols to show reactants and products.
• Balanced chemical equation : A balanced chemical equation has equal numbers of each type of atom on both sides of the arrow.
• Reactants : Reactants are the substances written on the left side of the arrow that take part in the reaction.
• Products : Products are the substances written on the right side of the arrow that form from the reaction.

📝 Essential Points

• Atoms are not created or destroyed in a chemical reaction, only rearranged.
• The mass of reactants equals the mass of products in any chemical reaction.
• In word and symbol equations, use a reaction arrow (à) rather than an equals sign because the two sides are different substances.
• A symbol equation is balanced by changing coefficients until the number of each kind of atom matches on both sides.
• In balanced equations, large numbers (coefficients) give the molecule ratio, while subscripts give the atom count within each molecule.
• Example: $2\mathrm{Cu}+\mathrm{O}_2\rightarrow 2\mathrm{CuO}$ balances copper and oxygen atoms on both sides.

💡 Memory Hook

Balance like a scale: coefficients adjust the load until every element matches on left and right.

📖 6. Group 1 metals and Group 7 halogens

🔑 Key Concepts & Definitions

• Group 1 metals : Group 1 metals are alkali metals that react in ways that become more vigorous as you move down the group.
• Group 7 halogens : Group 7 halogens are non-metals with similar properties and with reactivity that decreases as you move down the group.
• Universal indicator : Universal indicator is a lab dye used to monitor solutions during reactions by changing colour.

📝 Essential Points

• Group 1 metals burn in air with a characteristic flame colour and leave a white solid that turns universal indicator blue.
• In reactions with water, group 1 metals form fizz and a purple/blue indicator change, with severity increasing from lithium to potassium.
• Group 7 halogens exist as F2, Cl2, Br2, I2 as diatomic molecules and they form coloured vapours (yellow for fluorine, green for chlorine, red for bromine, purple for iodine).
• Halogen volatility and physical state follow the given boiling points: chlorine (−35 °C) and bromine (59 °C) and iodine (184 °C), so at 20 °C bromine is liquid and iodine is solid.
• Halogen reactivity decreases down the group: bromine is less reactive than chlorine, and more reactive halogens displace less reactive ones from their solutions.
• In Group 7 displacement, a more reactive halogen X2 pushes out a less reactive halogen Y2 from NaY to form Y2 and NaX, as shown by $X_2 + 2NaY \rightarrow Y_2 + 2NaX$.

💡 Memory Hook

Down Group 1 = more reactive; down Group 7 = less reactive (top of Group 7 displaces the bottom).

📖 7. Ionic tests and flame colours

🔑 Key Concepts & Definitions

• Flame colour test : A flame colour test identifies certain ions by the colour produced when a compound is heated in a flame.
• Silver nitrate test : A silver nitrate test identifies halide ions by reacting them with AgNO3 to form a solid product.
• Dilute nitric acid : Dilute nitric acid is added before AgNO3 so the halide ions react correctly during the silver nitrate test.

📝 Essential Points

• For the flame colour test, metal ions from Group 1 salts can give red (lithium), yellow-orange (sodium), or lilac (potassium) flames.
• For the silver nitrate test, add dilute nitric acid to the compound solution, then add drops of AgNO3 to detect halides.
• With chlorid e ions, the precipitation is white solid AgCl(s) formed by Ag+(aq)+Cl-(aq)→AgCl(s).
• The source notes halides detected by the silver nitrate test are Cl-, Br-, and I-.
• The flame colour test results are given alongside pH indicator turning blue for the Group 1 reactions with oxygen (as described).

💡 Memory Hook

AgNO3: Acid then drops—Cl- gives a white AgCl cloud/precipitate.

📖 8. Noble gases and their uses

🔑 Key Concepts & Definitions

• Noble gases : Noble gases are Group 0 elements that are very unreactive under normal conditions.
• Group 0 : Group 0 is the periodic-table group containing the noble gases.

📝 Essential Points

• Noble gases are described as inert because they do not readily form chemical reactions.

📊 Synthesis Tables

Group 1 vs Group 7 trends

Metals vs non-metals properties

⚠️ Common Pitfalls & Confusions

1. Mixing up atomic number with mass number: atomic number = protons (and electrons in a neutral atom), while mass number = protons + neutrons.
2. For ions, forgetting charge comes from electron change: losing electrons makes a positive ion, gaining electrons makes a negative ion
3. Using isotope idea wrongly: isotopes change number of neutrons (same atomic number), so you shouldn’t think mass number of a specific isotope must be memorised.
4. Writing ionic formulas without balancing total charges to neutrality, or forgetting Roman numerals for metals that form multiple positive ions.
5. Balancing equations by changing subscripts instead of only changing coefficients, or forgetting atoms must match on both sides.
6. Confusing Ar calculation with mass number: Ar is the weighted average using isotope abundances; mass number is just proton + neutron.
7. Assuming group trends are about individual properties only rather than patterns: e.g., Group 1 reactivity increases down the group; Group 7 decreases down the group.

✅ Exam Checklist

1. Define atomic nucleus, atomic number, mass number, and isotopes using the relationships given.
2. Explain why atoms have no overall charge, and describe how simple ions (e.g. Na+, Mg2+, Cl–, O2–) form from gaining/losing electrons.
3. Given an atom/ion, use atomic number and mass number to find numbers of protons, neutrons, and electrons (including simple ions).
4. Use isotopic abundances to calculate Ar via the required weighted average method (including the chlorine example idea).
5. Describe periodic table layout: elements increase in atomic number left-to-right and top-to-bottom, with groups and periods and the link to outer-shell electrons and occupied shells.
6. State and use the electron-shell capacity pattern (2 in first shell, 8 in second, then 8 then 18) to determine electron configurations for elements up to the course scope.
7. Use group membership (Group 1 and Group 7) to predict reactivity trends and the stated boiling/melting point trends.
8. Write correct ionic compound formulae by selecting ion symbols/charges, then balancing so total positive = total negative (using Roman numerals when needed).
9. Calculate Mr from a molecular formula (using Ar values) and from a compound formula with multipliers (e.g. Na2CO3).
10. Write, interpret, and balance chemical equations: identify reactants/products, use à (not equals), and balance with coefficients so each element has equal atom counts.
11. For Group 1 metals: recall and explain the characteristic observations with oxygen/air, water, and the given flame colours/indicator change patterns.
12. For halogens: recall diatomic molecules (F2, Cl2, Br2, I2), colour of vapours, volatility/physical states using the given boiling points, and displacement rule with the equation format X2 + 2NaY → Y2 + 2NaX.
13. Perform flame colour and silver nitrate tests: match Li+, Na+, K+ flame colours, and for halides use dilute nitric acid then AgNO3 to produce white/cream/yellow precipitates (AgCl/AgBr/AgI) and write the ionic equation.
14. State uses and inert nature for noble gases (Group 0): helium, neon, and argon (at least one use each) and why they are inert.