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CHEM1110 > Week 1 > Flashcards

Week 1 Flashcards

1
Q

Define chemistry

A

The characterisation, composition and transformation of matter

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2
Q

Components of the scientific method

A
  • Observations
  • hypothesis
  • experimentation
  • theory, scientific law
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3
Q

Scientific theory

A

A well established hypothesis

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4
Q

Scientific law

A

Has no known exceptions, regarded as fact in modern science

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5
Q

International unit system

A

Metric system based on the SI units

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6
Q

What are the SI units

A
  • kilogram (mass)
  • metre (length)
  • second (time)
  • kelvin (temp)
  • mole (amount of a substance)
  • ampere (electric current)
  • candela (luminous intensity)
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7
Q

Mass

A

measure f amount of matter in a body. Identical irrespective of location

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8
Q

Weight

A

Function of gravitational attraction and therefore varies

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9
Q

Litre

A

1 decimeter cubed

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10
Q

Angstrom

A

0.1nm (10 to the power of -10m), about the size of an atom

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11
Q

Density

A

mass/volume (SI unit is kg/m3)

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12
Q

kelvin

A
  • temperature scale used in chemistry

- defined by absolute zero (temp at which molecular motion ceases)

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13
Q

Celcius to Kelvin

A

K = C* + 273

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14
Q

Basics of scientific notation

A

Numbers expressed as a number between 1 and 10 multiplied by 10 with a whole number power (exponential power)

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15
Q

Precision

A

how close your results are to the true results

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16
Q

Accuracy

A

how close the results are to each other

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17
Q

Rules for significant figures

A
  • any non-zero digit is significant
  • zeros that proceed the first non-zero digit are NOT significant
  • zeros between non-zero digits are always significant
  • zeros trailing non-zero digits are only significant if the number contains a decimal point
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18
Q

Leading zeros (significant numbers)

A

Any zeros before a non-zero digit

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19
Q

Adding/subtracting significant numbers

A

Significant figures to the right of the decimal point in the final answer are determined by the lowest number of significant figures to the right of the decimal point in the added numbers. Add together then round to that many significant numbers.

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20
Q

Multiplying/dividing significant figures

A

Significant figures in the product are determined by original number possessing the least number of significant figures

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21
Q

Matter

A

Any substance that occupies space and possesses mass

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22
Q

Primary states of matter

A
  • solid: rigid, fixed shape/volume, incompressible
  • liquid: fluid, fixed volume but no shape, incompressible
  • gasses: fluid, no fixed volume/shape, compressible
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23
Q

Vapour

A
  • gaseous matter at a liquid or solid boundary (in equilibrium with a solid or liquid)
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24
Q

Plasma

A

An electrically neutral gas of ions and electrons

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25
Change in matter
- physical change: change in form but not chemical identity | - chemical change: transformed into a new kind of matter
26
Physical properties of matter
characteristics observed for a material without changing chemical identity. - melting/boiling point, colour, taste, odour
27
Chemical properties of matter
characteristics observed for a material that involves chemical change - ability to lose or gain electrons, electronegativity
28
Classification of matter
Matter -----> Mixtures or Substances Mixtures ---> heterogeneous or homogeneous Substances ---> compounds or elements
29
Mixtures
combination of two or more different substances that are not combined mechanically. Have physical means of separation.
30
Sub-classes of mixtures
heterogenous and homogeneous
31
heterogeneous mixtures
Mixtures with physically distinct components (e.g. soil, sand)
32
Homogeneous mixtures
physically uniform (e.g. solutions, air)
33
Pure substances
Can not be broken down further by physical processes. They have a uniform appearance and have a fixed definite composition. They possess distinct physical chemical properties and are the same around the world. Can be divided into elements and compounds.
34
Number of naturally occurring elements
92
35
Liquid elements at room temp
bromine and mercury
36
The majority of elements are...
metals
37
Compounds
composed of 2 or more elements that combine in fixed and definite proportions by mass, most pure substances are compounds. They can be decomposed by chemical means to two or more elements.
38
How to elements in compounds combine together?
Elements in compounds combine together in simple whole number ratios
39
Molecular compounds vs ionic compounds
What is formed in a molecular compound is a discrete unit (i.e. one molecule of whatever is formed) Ionic compounds do not form discrete units because they are charged ions and move around
40
Law of conservation of mass
- no loss or gain in mass occurs in chemical processes | - Lavoisier's (french scientist, late 1700s) experiments showed mass of products matched that of the reaction
41
Law of definite proportions
- Elements in a given chemical compound always combine in the same proportions by mass - Joseph Proust: compounds prepared in the lab were identical to those isolated from natural sources
42
Law of multiple proportions
- When two elements can form more than one compound, the different masses of one element combining with the same mass of the second element will be in a simple whole number ratio - john dalton carbon and oxygen experiment
43
Dalton's conclusions
- elements made up of tiny particles (atoms) - atoms of a given element were identical but different to those of another element - compounds: product of atoms of different elements combining in specific ways in fixed ratios - chemical reactions involved reorganisation of atoms and the ways they were bound together but did not themselves undergo change
44
Joseph Gay Lussac
- corrected Dalton's formulations for compounds using volumes of gases
45
Thompson and the electron
Thompson determined the electrical charge to mass of the electron ratio (e/m = - 1.76 x 10 *8 coulombs per gram)
46
Millikan
- used thompsons work of e/m to deduce the charge of an electron (1.6 x 10 * -19 C) and the mass of the electron (9.1 x 10* -28 g)
47
Features of the atom apparent by 1900s
- contained negatively charged electrons | - electrically neutral overall
48
Rutherford's gold foil conclusions
- the massive deflections result of interactions between like charges - nucleus contained highly concentrated cor of positive charge (protons) - later experiments showed protons contain an equal but opposite charge to an electron possessing mass of 1.6 x 10* -27 (1800x mass of electron)
49
relative size of subatomic particles
nucleus occupies 1x 10* -15% of total atom volume atomic radius: 10* -10m Nuclear radius: 5x10* -15
50
Neutron
fundamental subatomic particle with a mass almost identical to protons. Carries no charge
51
Proton mass and charge
Mass: 1.67 x 10* -27 Charge: 1.6022 x 10* -19 (approx 1)
52
Neutron mass and charge
Mass: 1.67 x 10* -27 Charge: 0
53
Electron mass and charge
Mass: 9.11 x 10* -31 Charge: -1.6022 x 10* -19 (approx -1)
54
Atomic number
- the number of protons in an atom - given symbol Z - characterises an individual atom
55
Mass number
- sum of protons and neutrons present in the nucleus of an atom - symbol A - measured in unified atomic mass units (u) or Daltons
56
1u (Da)
1u = 1Da | = 1.660538782 x 10* -27 kg
57
How to calculate # of neutrons in an atom
A - Z = # neutrons | mass number - atomic number
58
Isotopes
- atoms of the same elements with different numbers of neutrons (and therefore mass numbers)
59
Average atomic mass
- typically not whole number - average mass of atoms and isotopes of an element - shown on periodic table - calculations to do with average atomic
60
Calculating average mass
average mass = mass1 x abundance + mass2 x abundance2......
61
Allotropes
elements where the atoms connect together in more than one way to form chemically different substances (can have different physical characteristics)

CHEM1110 (5 decks)

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