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CHEMISTRY > amount of substance > Flashcards

amount of substance Flashcards

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1
Q

what is the symbol for amount of substance ?

A

n

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2
Q

what is the unit used to measure amount of substance ?

A
  • mole
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3
Q

what does avogadro constant represent ?

A
  • the number of atoms per mole of the carbon-12 isotope.
  • 6.02 x 10^23
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4
Q

how to calculate the mass of 1 mole of the element ?

A
  • one mole of a substance will have same mass in grams as the relative atomic mass or formula mass of that substance
  • mass of 1 mole = relative atomic mass in grams
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5
Q

how to calculate moles when mass and molar mass are given?

A
  • moles (mol) = mass (g) / molar mass (g mol^-1)
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6
Q

define relative atomic mass (Ar)

A
  • the average mass of an atom of an element compared with 1/12 the mass of one atom of carbon-12
  • Ar = average mass of an atom of an element
    ——————————————————
    1/12 the mass of one atom of carbon-12
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7
Q

Define Relative molecular Mass. (Mr)

A
  • the sum of the relative atomic masses of all the atoms in a molecule.
  • Mr = average mass of one molecule of an element of compound
    —————————————————————
    1/12 the mass of one atom of carbon-12
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8
Q

define avogadro constant

A
  • number of particles / atoms / ions in one mole of a substance
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9
Q

Helium is a monatomic gas. He. How many atoms are in 0.200 mol helium gas ?

A
  • step 1: 1 mol helium contains 6.023 x 10^23 atoms
  • step 2: 0.200 mol helium contains 0.2 x 6.023 x 10^23 = 1.205 x 10^23 atoms.
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10
Q

what is the mass in grams of 2.5 moles of ethanol, C2H5OH?

A
  • step 1: Mr ethanol = (2x12) + (5x1) + 16 + 1 = 46
  • step 2: Mass (g) 2.5 x 46 = 115 g
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11
Q

State Avogadros law.

A
  • under the same temp and pressure - one mole of any gas would occupy the same volume .
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12
Q

how much volume does a gas occupy at room temp and pressure ?

A
  • 24dm^3 // 24000 cm^3
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13
Q

define molar gas volume

A
  • the volume per mile of gas molecules
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14
Q

what are the conversions

A

dm^3 —> cm^3 (x1000)
cm^3 —> dm^3 (divide1000)

dm^3 —> m^3 (divide1000)
m^3 —> dm^3 (x1000)

cm^3 —> m^3 (divide 1000000)
m^3 —> cm^3 (x1000000)

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15
Q

State Boyles law

A
  • at constant temp (T), the volume V of a fixed mass of gas is inversely proportional to the pressure p applied to it
  • p x V = constant
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16
Q

State Charles law

A
  • at constant pressure, the volume of a fixed mass of gas is proportional to its temperature .
  • V = constant x T
17
Q

what are the assumed properties of an ideal gas ?

A
  • made up of identical particles in continuous random motion.
  • parties thought of as point-like - position but with zero volume (volume of gas particle taken to be 0)
  • particles don’t react when they collide.
  • collisions between particles perfectly elastic - total kinetic energy of particles after collision is same as before collision
  • particles have no intermolecular forces - don’t attract/ repel each other .
18
Q

why do different gas occupy the same volume?

A
  • the gas particles are very spread out, hence individual differences has no effect .
19
Q

how to calculate moles when gas volume is given ?

A
  • moles (moles) = volume (dm^3) / 24

or

  • moles (mol) = volume (cm^3) / 24000
20
Q

worked example = 3.25 g zinc reacted with an excess of HCl to form zinc chloride + hydrogen gas , calculate the volume of hydrogen has (in cm^3) formed in this reaction

Zn + 2HCl —> ZnCl2 + H2

A
  • use the mass of zinc to work out moles
  • moles of zinc = 3.25/65 = 0.05 mol
  • there’s a 1:1 ratio between zinc + hydrogen (of 0.05 mol of zinc reacting , 0.05 moles hydrogen formed)
  • volume of hydrogen = moles x 24
  • 0.05 x 24 = 1.2dm^3
  • 1.2 x 1000 = 1200 cm^3
21
Q

write down the ideal gas equation (in words + symbols including symbols )

A
  • pV = nRT
  • pressure (Pa) x volume (m^3) = number of moles (mol) x gas constant (8.314 Jmol-1) x temp (K)
22
Q

1 atmospheric pressure is equal to how many pascal?

A
  • 1 atm = 101325 Pa
23
Q

how do you convert between K and C temperatures ?

A
  • degrees to K = +273
  • k to degrees = -273
24
Q

define empirical formula

A
  • the simplest whole number ratio of atoms of each elements in a compound .
25
Q

describe what’s meant by chemical formula .

A
  • the actual number of atoms of each element used to make a molecule of formula unit .
26
Q

define percentage yield .

A
  • the % of a product produced by a reaction, compared to a theoretical maximum .
27
Q

how would you calculate percentage yield ?

A
  • mass of actual yield / theoretical yield ( x100 )
28
Q

what can the percentage yield of a practical be used to investigate ?

A
  • efficiency of practical techniques and whether reactions proceed as estimated
29
Q

define atom economy .

A
  • % of amount of reactants made into a certain (useful) product.
30
Q

how would you calculate atom economy ?

A
  • mr of desired product / sum of molecular masses (mr) for all reactants (x100)
31
Q

what can the atom economy of a reaction be used to investigate ?

A
  • efficiency of using a specific reaction to produce a product
32
Q

what can the atom economy of a reaction be used to investigate ?

A
  • efficiency of using a specific reaction to produce a product
33
Q

define ionic equations

A
  • concise method of writing down the important changes to the ions in a chemical reaction .
34
Q

consider how to work out the ionic equation for the reaction between hydrochloric acid and sodium hydroxide that produces sodium chloride and water

A
  • write full equation = HCl (aq) + NaOH (aq) —> NaCl (aq) + H2O
  • write equation as ions and cancel the ions that appear on both sides:
  • HCl (aq) + Cl- (aq) + Na+ (aq) + OH- (aq) —> Na+ (aq) + Cl- (aq) + H2O (l)
  • ionic equation = H+ (aq) + OH- (aq) —> H2O (l)
35
Q

define spectator ions .

A
  • ions that appear unchanged on both sides of the ionic equation are cancelled out
36
Q

what’s the equation that links mol, concentration and volume ?

A

moles = conc x volume

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