Myanmar Matriculation Chemistry flashcards

33 cards. Read the term and the definition here. The deck at the top of this page is optional.

Mole definition

One mole is the amount of substance containing 6.022 × 10²³ elementary particles, equal to the Avogadro constant. Its mass in grams equals the relative formula mass of the substance.

Molar volume at STP

One mole of any gas occupies 22.4 dm³ (22 400 cm³) at standard temperature and pressure, 0 °C and 1 atm. Use n = V/22.4 with V in dm³.

Empirical formula

The simplest whole-number ratio of atoms of each element in a compound. Find it by dividing each percentage by atomic mass, then dividing all results by the smallest quotient.

Limiting reactant

The reactant that is completely used up first and so fixes the maximum amount of product. Identify it by dividing each reactant's moles by its coefficient in the balanced equation.

Percentage yield

Actual mass of product obtained divided by theoretical mass predicted from stoichiometry, multiplied by 100. Values below 100 arise from side reactions, incomplete reaction and handling losses.

Dilution equation

For dilution with solvent, C₁V₁ = C₂V₂ because the number of moles of solute is unchanged. Concentrations must use the same unit and volumes the same unit.

Isotopes

Atoms of the same element with the same proton number but different neutron numbers, hence different nucleon numbers. They have identical chemical properties but slightly different physical properties and masses.

Ionisation energy

The energy needed to remove one electron from each atom in one mole of gaseous atoms, forming singly positive gaseous ions. It increases across a period and decreases down a group.

Electronegativity

The tendency of an atom in a covalent bond to attract the shared pair of electrons toward itself. It increases across a period and decreases down a group.

Ionic bonding

Electrostatic attraction between oppositely charged ions formed by electron transfer, typically from a metal to a non-metal. Compounds have high melting points and conduct when molten or in aqueous solution.

Hydrogen bonding

An unusually strong intermolecular attraction between a hydrogen atom bonded to N, O or F and a lone pair on a nearby N, O or F atom. It explains water's high boiling point and ice's low density.

Oxidation number

A number showing the charge an atom would have if bonding electrons were assigned to the more electronegative atom. Free elements are 0 and the sum in a neutral compound is 0.

Oxidising agent

A species that accepts electrons from another substance, thereby oxidising it, and is itself reduced. Its oxidation number falls during the reaction.

Electrolysis electrodes

At the negative cathode, cations gain electrons and are reduced. At the positive anode, anions lose electrons and are oxidised. Electrons flow through the external circuit, not the electrolyte.

Faraday's law of electrolysis

The mass of substance deposited at an electrode is proportional to the quantity of charge passed: m = (Q × M)/(z × F), where z is the number of electrons per ion and F is 96 500 C mol⁻¹.

Catalyst effect

A catalyst provides an alternative reaction pathway of lower activation energy, increasing the rate of both forward and reverse reactions. It is not consumed and does not change the enthalpy change.

pH definition

pH = −log₁₀[H⁺], where [H⁺] is hydrogen ion concentration in mol dm⁻³. A change of one pH unit means a tenfold change in hydrogen ion concentration.

Strong versus weak acid

A strong acid ionises almost completely in water, giving a high [H⁺] and low pH. A weak acid ionises only partially, establishing an equilibrium, so its pH is higher at the same concentration.

Neutralisation reaction

The reaction of an acid with a base to form a salt and water. With a carbonate it also produces carbon dioxide; with a reactive metal it produces hydrogen instead of water.

Solubility rule for sulfates

Most sulfates are soluble in water, but barium sulfate, lead(II) sulfate and calcium sulfate are insoluble or only slightly soluble. Acidified barium chloride therefore tests for sulfate ions.

Flame test colours

Sodium gives yellow, potassium lilac, lithium crimson, calcium brick-red, copper blue-green and barium apple-green. A clean nichrome wire and concentrated acid are used.

Test for carbon dioxide

Bubbling the gas through limewater, calcium hydroxide solution, turns it milky or cloudy because insoluble calcium carbonate is precipitated. Excess carbon dioxide redissolves the precipitate.

Alkene test

Bromine water is decolourised from orange-brown to colourless by addition across the carbon–carbon double bond. Alkanes give no such reaction in the dark, so the test distinguishes them.

Esterification

The reversible reaction of a carboxylic acid with an alcohol, catalysed by concentrated sulfuric acid, producing an ester and water. The ester usually has a characteristic sweet fruity smell.

Exothermic reaction

A reaction that releases heat to the surroundings, so the temperature rises and ΔH is negative. The products have less enthalpy than the reactants because bond formation releases more energy than bond breaking absorbs.

Hess's law

The enthalpy change of a reaction depends only on the initial and final states, not the route. It allows ΔH to be calculated from formation enthalpies or bond energies.

Calorimetry formula

Heat released or absorbed is q = mcΔT, where m is the mass of water in grams, c is 4.2 J g⁻¹ K⁻¹ and ΔT is the temperature change in kelvin or Celsius degrees.

Le Chatelier's principle

When a system at equilibrium is disturbed by a change in concentration, pressure or temperature, the position of equilibrium shifts so as to oppose that change. Catalysts do not shift it.

Equilibrium constant Kc

For a balanced reversible reaction, Kc equals the product of product concentrations raised to their coefficients divided by the same for reactants. It varies only with temperature.

Standard electrode potential

The potential of a half-cell measured against the standard hydrogen electrode, which is assigned 0.00 V, under standard conditions of 298 K, 1 atm and 1 mol dm⁻³.

Galvanic cell polarity

In a galvanic cell the anode is the more reactive metal and is the negative terminal, where oxidation occurs. The cathode is the positive terminal, where reduction occurs.

Test for ammonium ion

Warming the salt with sodium hydroxide solution releases ammonia gas, which turns damp red litmus paper blue. The reaction is NH₄⁺ + OH⁻ → NH₃ + H₂O.

Relative atomic mass

The weighted average mass of the naturally occurring isotopes of an element compared with one-twelfth the mass of a carbon-12 atom. It is not a whole number for most elements.