ATOMIC STRUCTURE

2. ATOMIC STRUCTURE

3. ATOMIC MODEL Proton, p (positive) Neutron, n (neutral) Electron, e (negative) Nucleons as the atomic nucleus. Moves around the atomic nucleus • Points to ponder: • All the elements, except hydrogen atom, contain neutrons in their atomic nuclei. • Distinguish between nucleus, neutron and nucleon??

4. ATOMIC SYMBOL(ISOTOPE NOTATION) • Atomic Symbols: • The atom of each element is made up of electrons, protons and neutrons. All atoms of the same neutral element have the same number of protons and electrons but the number of neutrons can differ. Atoms of the same element but different neutrons are called isotopes. Because of these isotopes it becomes necessary to develop a notation to distinguish one isotope from another - the atomic symbol. The atomic symbol has three parts to it: • 1. The symbol X: the usual element symbol • 2. The atomic number A: equal to the number of protons (placed as a left subscript) • 3. The mass number Z: equal to the number of protons and neutrons in the isotope (placed as a left superscript)

5. We also can calculate the mass by using : • Relative atomic mass A (X) = average mass of 1 atom X 1/12 (X )mass of 1 atom C-12 • Relative molecular mass( molecular weight) M (X) = average mass of 1 molecule X 1/12 (X) mass of 1 atom C-12 Important: both calculation does not have UNITS.

6. MOLE CONCEPT • MOLE, n : The amount of substance thaht contains as many elementary aprticles as there are atoms in exactly 12.000g of carbon-12. I mol = 6.02×10²³ particles (atoms mlecules) • MOLAR MASS, M : The mass of one mole of a substance molar mass, M = Ar or Mr (g/mol)

7. Mole, n = number of particles Avogadro constant, NA 1 • 1 mol = 6.02×10²³ particles ( Avogadro constant = 6.02×10²³ particles molˉ¹ ) • 1 mol = Ar or Mrin unit gram Ar of an element is obtained from the Periodic Table Mr of compound = ∑ Ar of each constituent atom The mass must be expressed in unit gram Mole, n = mass, m molar mass, M 2

8. Mole, n = volume of gas molar volume • I mol of any gas = 22.4 L at s.t.p (P = 1 atm ; T = 273K ) = 24.5 L at SLC (std. lab conditions) (P = 1atm ; T = 298K ) • 1 mol ≠ 1molecule • Physical quantity:unit: Atomic/molecular mass u (atomic mass unit) Molar mass g/mol Ar or Mr & molecular weight DIMENSIONLESS Mass of substance g 3

9. CONCENTRATION OF MEASUREMENTS • Amount concentration (molarity) c = mole solute, n volume solution, v unit: mol Lˉ¹, mol dmˉ³, M • Molality m = mole solute mass solvent ,kg unit: mol kgˉ¹, molal, m • Mole fraction Xa = numbers of moles of components in the mixture number of moles of unit: X all substance in the mixture

10. Per cent by mass % w/w = mass solute × 100% mass solution * mass solution= mass solute + mass solvent • Per cent by volume % v/v = volume solute × 100% volume solution * volume solution= mass of solution density of solution

11. OXIDATION NUMBER • DEFINITION: - For monatomic ions, the oxidation number is equal to the charge on the ion. - For covalently bonded atoms, the oxidation number is the charge on an atom calculated by assigning both electrons of a shared pair to the more electronegative atom. • RULES: - In free elements(that is an element not combined chemically with different element) each atom has a oxidation number of zero(0) (H2, Br2, Na, O2) - For ions composed of only one atom (that is monatomic ions) the oxidation number is equal to the charge on the ion. (S²ˉ= -2) - In a neutral molecule, the sum of the oxidation numbers of all atoms must be zero (CO2, NaCl) - In a polyatomic ion, the sum of oxidation numbers of all the elements in the ion must be equal to the net charge of the ion. (NO3ˉ)

12. REDOX REACTION • Definition: Reactions that involve both reduction and oxidation are called redox reactions - Oxidation : Removal (loss) of electrons species will get less negative or more positive - Reduction : Gain of electrons species will become more negative or less positive Example: H + Cl2 The hydrogen is oxidized(increased in oxidation number) and the chlorine is reduced (decreased in oxidation number). Thus, the reducing agent is the H2 and the oxidizing agent is Cl2. 2HCl

13. CHEMICAL EQUATION • A chemical equation is a way of denoting a chemical reaction using the symbols (chemical formulae) for the participating particles (atoms, molecules, ions and free radicals). CaCO3 (s) + 2HCl (aq)CaCl2 (aq) + CO2 (g) + H2O (l) • A reactant is a starting substance in a chemical reaction (on the left) • A product is a substance that results from a reaction (on the right) *One can also indicate the conditions under which a reaction takes place, as well as the presence of a catalyst.

14. BALANCING A CHEMICAL EQUATION (REDOX) • General rules: • Write a correct formulae for all reactants and products on the correct side of the “reaction arrow” • The equation can be balanced only by adjusting the coefficients of the formula, as necessary • Never introduce extraneous formulae (spectator ions) that are not involved in the reaction • The total numbers of the atoms of the elements is the same on both side. Examples of balancing redox reactions

15. ION-ELECTRON METHOD • Separate the equation in two half-reactions(acidic and basic) • Balance the atoms of the elements undergoing changes in oxidation number • Balance oxygen atoms by using H2O • Balance hydrogen atoms by using H+ • Balance the charge in each half-reaction by adding electrons to equalize the ionic charges • Multiply the half-reactions by appropriate integer to ensure that the number of electrons lost in oxidation is equal to the number of electrons gained in reduction • Sum the half-reactions. Simplify the overall equation algebraically Examples

16. The End