Chemistry > Unit 1 > Flashcards

Unit 1 Flashcards

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1
Q

The particulate nature of matter means that…

A

All matter is made up of particles (matter is particulate).

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2
Q

Matter

A

Something that has mass and occupies space. Can be classified as pure substance or mixture.

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3
Q

Types of matter

A

Elements, compounds, mixtures.
Elements and compounds are pure substances.

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4
Q

Elements

A

Primary constituents of matter, which cannot be further broken down chemically (by chemical methods). Elements are fundamental components and have one type of atom

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5
Q

Compounds

A

Consists of more than one type of atom (atoms of different elements) chemically bonded together in a fixed ratio. The property of a compound is different from its constituent elements.

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6
Q

Mixtures

A

Contain more than one element and/or compound in no fixed ratio. Not chemically bonded and so can be separated by physical methods. Retains the properties of its components.

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7
Q

Mixture can be classified as…

A

homogeneous or heterogeneous

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8
Q

Homogeneous mixture

A

Mixture with uniform and definite composition

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9
Q

Heterogeneous mixture

A

Mixture with non-uniform composition

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10
Q

Sublimation

A

solid to gas

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11
Q

Deposition

A

Gas to solid

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12
Q

Condensation

A

Gas to liquid

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13
Q

Gas to plasma

A

Ionisation

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14
Q

Plasma to gas

A

De-ionisation

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15
Q

Celsius to Kelvin

A

Celsius + 273.15K

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16
Q

Nuclear symbol notation

A

Upper number is mass number (A)
Lower number is atomic number (Z)
Chemical symbol for the element in the middle

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17
Q

Isotopes

A

Atoms of the same element with different numbers of neutrons.
Same chemical properties, different physical properties.

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18
Q

Why do isotopes have the same chemical properties?

A

Because chemical behaviour is determined by the number of electrons and isotopes have the same number of electrons, and neutrons don’t take part in electronic configuration.

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19
Q

Why do isotopes have different physical properties?

A

Because they have a different mass since they have different numbers of neutrons. Physical properties are dependent on atomic mass.

Heavier isotopes have lower zero-point energy, so more energy is required to break the bonds and dissociate the molecules.

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20
Q

Relative atomic mass

A

the ratio of the average mass of atoms of an element to 1/12th of the mass of an atom of Carbon-12.

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21
Q

The RAM depends on…

A

the abundance of different isotopes in nature.

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22
Q

RAM formula

A

sum(mass x abundance)/total abundance

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23
Q

Mass spectrometer

A

An analytical device used to determine the abundance of various isotopes of an element based on a measurement of the m/z ratio of ions within a sample of the element.

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24
Q

mass-to-charge (m/z) ratio

A

The ratio of the mass of an ion divided by its charge.

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25
Stages of the mass spectrometer
Vaporisation into gaseous molecules. Ionisation to form ions. Acceleration into an electric field. Deflection by an electromagnet; degree of deflection depends on m/z ratio (low m/z ratio = high deflection, hhigh m/z ratio = low deflection) Detection - gaseous ions are detected and a mass spectrum is produced.
26
Mass spectrum
A representation of the m/z ratio versus abundance (relative).
27
Spectroscopy
The study of the interaction between matter and light. Elements can be observed using spectroscopy by observing their unique emission spectra.
28
Emission spectra (emission line spectrum)- what they are and how they are produced.
the range of frequencies or wavelengths of electromagnetic radiation emitted during an electron transition from a higher to a lower energy level. Emission spectra are produced by atoms emitting protons when electrons in excited state return to lower energy levels or ground state. Emission line spectrum is produced when energy is applied to a sample of gas then viewed through a spectroscope.
29
Continuous spectrum
A spectrum containing all the frequencies or wavelengths across a range of electromagnetic radiation.
30
Spectroscope
A device for viewing the emission spectra of elements in the visible region.
31
Line in the emission spectrum converge at...
High energy.
32
Each element produces its unique emission line spectrum, what is one application of the emission line spectrum.
To identify unknown elements.
33
Principal quantum number
n - the integer number given to a main energy level.
34
Ground state
Lowest energy state of an atom when all its electrons occupy their lowest energy levels.
35
Electron transition
The movement of an electron between the energy levels in an atom accompanied by the absorption or emission of energy.
36
In what form is energy absorbed or emitted?
Photons (small packets of energy).
37
Principal/main energy levels
The discrete energy levels in which the electrons are located.
38
How many electrons can a main energy level hold?
2n^2 electrons.
39
Heisenberg uncertainty principle
A scientific principle that states we cannot know both the exact location and velocity of an electron at the same time.
40
Main energy levels are divided into sublevels. What are the letters assigned to sublevels?
s, p, d, f - based on the shape of the atomic orbitals they contain.
41
Atomic orbital
A region of space where there is a high probability of finding an electron.
42
How many electrons can each atomic orbital hold?
Maximum of 2.
43
How many atomic orbitals do each sublevel consist of?
s: 1 p: 3 d: 5 f: 7
44
Electron configuration
The arrangement of electrons in their different levels around the nucleus of an atom.
45
The Aufbau principle
Electrons fill the atomic orbital of the lowest energy levels first.
46
The Pauli exclusion principle
An atomic orbital can only hold 2 electrons and they must have opposite spins (spin up or spin down).
47
Hund's rule
Electrons fill degenerate orbitals singly before doubling up.
48
Degenerate orbitals
Orbitals within the same sublevel of an atom that have the equal energy levels.
49
Exceptions to the Aufbau principle
Copper and chromium.
50
Orbital diagrams/Orbital notations represent...
how electrons occupy the atomic orbitals of an atom.
51
What force holds the protons and neutrons together?
Strong nuclear force.
52
By what force are electrons held in position
Electrostatic attraction between protons and electrons.
53
Ionisation energy
The minimum amount of energy required to remove one mole of electrons from one mole of gaseous atoms or ions to form one mole of gaseous 1+ or other ions.
54
Ionisation
The process of removing electrons from an atom in its ground state (transition to n = infinity).
55
What corresponds to ionisation in an emission spectrum?
The limit of convergence at higher frequencies.
56
Factors affecting the strength of electrostatic attraction
Nuclear charge - the number of protons in the nucleus Atomic radius - distance between the nucleus and the outer electrons Shielding effect - repulsion between electrons in the inner shell and valence electrons in the outer shell.
57
What is the general trend in ionisation energy within periods?
1st ionisation energy increases within a period of elements (increasing atomic number order) due to increased nuclear charge and attraction between electrons and nucleus leading to a decreased atomic radius, while the shielding remains relatively constant. There are a few exceptions (e.g. Be to B and N to O - explain why).
58
Group 1 vs Group 7 elements IEs
Group 1 elements have low ionisation energies, while noble gases have high ionisation energies. Explain how noble gases have lower mps/bps than group 1 elements even though they have higher ionisation energies.
59
Convergence limit
The point at which the energy levels/spectral lines in an atom converge (frequency or wavelength that corresponds to ionisation).
60
Energy formula
E=hf Where: E = energy (J) h = Planck's constant (6.63 x 10^-34) (Js) f - frequency (s^-1)
61
Speed of light formula
c = fλ Where: c = speed of light (3.00 x 10^8) (ms^-1) f = frequency (s^-1) λ = wavelength (m)
62
Successive ionisation energies
The minimum energy required to remove each electron from an ion in a gaseous form that is becoming increasingly positive.
63
Successive ionisation
The process of removing successive electrons from a gaseous atom or positive ion.
64
Why do successive ionisation energies increase?
The electrons are being removed from an increasingly positive ion. This leads to an increase in electrostatic attraction between the nucleus and remaining electrons, which incraeses ionisation energy.
65
Why is the increase in ionisation energy not constant?
There is a larger increase where the main energy level in which the eletron is removed from changes (this is further evidence for the existence of main energy levels).
66
What is the application of the successive ionisation energies?
It can be used to determine which group the element beloongs to, by counting the number of electrons ionised before the first large increase. (This is the number of electrons in the valence shell.)
67
Molar mass units
g mol^-1
68
Empirical formula of a compound
The simplest ratio of atoms of each element present in that compound.
Chemistry
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