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chemistry > 1: Atomic Structure and the Periodic Table > Flashcards

1: Atomic Structure and the Periodic Table Flashcards

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1
Q

What is the basic structure of an atom

A
  • Empty space around a small dense nucleus
  • Nucleus contains protons and neutrons
  • Nuclei of atoms have an overall positive charge
  • Electrons found in orbitals around the nucleus
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2
Q

What is relative mass and charge of a proton

A
  • Charge: +1
  • Mass: 1
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3
Q

What is the relative mass and charge of a neutron

A
  • Charge: 0
  • Mass: 1
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4
Q

What is the relative mass and charge of an electron

A
  • Charge: -1
  • Mass: 1/1836 (negligible)
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5
Q

What is atomic number (proton number) of an element

A

The number of protons in the nucleus of an atom, which is equal to the number of electrons in a neutral atom of an element

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6
Q

What is the mass number (nucleon number) of an element

A

The total number of protons and neutrons in the nucleus of an atom

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7
Q

How to calculate number of neutrons

A

Number of neutrons = mass number - atomic number

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8
Q

What are ions

A

Formed when atoms (which are neutral) gain or lose an electron, causing them to be charged

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9
Q

What is an isotope

A

Atoms of the same element with the same number of protons and electrons but a different number of neutrons, displaying the same chemical characteristics but different physical ones

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10
Q

Why do isotope have different physical properties to other isotopes of the same element

A
  • They have different numbers of neutrons
  • Since neutrons are neutrally charged, they only add mass
  • Causing the isotopes to have differences is mass, density, melting, and boiling points
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11
Q

What is the symbol of an isotope

A

The chemical symbol/word followed by a dash and the mass number (e.g. Carbon-12)

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12
Q

How is relative atomic mass calculated, and its equation

A

From the mass number and relative abundances of all the isotopes of a particular element

(Relative abundance i1 x mass i1) + (relative abundance i2 x mass i2) …/100

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13
Q

Define relative isotopic mass

A

The mass of an isotope relative to 1/12 of a carbon-12 atom

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14
Q

Define relative atomic mass (Ar)

A

The weighted mean mass of an atom relative to 1/12 of the mass of a carbon-12 atom

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15
Q

Define relative formula mass (Mr)

A

The total mass of the substance, calculated by adding up the relative atomic masses of all the atoms present in the formula

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16
Q

What is mass spectra

A
  • Peaks show abundance against mass for elements with isotopes
  • It can also show fragments of molecules (e.g. CH3+)
  • It creates fragments by bombarding the molecule with electrons to create molecular ions
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17
Q

What is the M+ peak in mass spectra

A
  • Shows the more abundant isotope
  • Is the molecular ion peak for molecules (has the highest m/z value and is equal to the relative molecular mass of the compound)
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18
Q

What is the M+1 peak in mass spectra

A

The smaller peak in molecular mass spectra that is due to the natural abundance of the isotope carbon-13

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19
Q

What is the mass spectra of chlorine and how is the ratio determined

A
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20
Q

Define ionisation energy (IE)

A

The amount of energy required to remove one mole of electrons from one mole of gaseous atoms of an element to form one mole of gaseous ions, under standard conditions

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21
Q

What are the units of ionisation energy

A

Kilojoules per mole

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22
Q

Define the first ionisation energy (IE1), and write the equation of IE1 for calcium

A

The energy required to remove one mole of electrons from one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions, under standard conditions

23
Q

Define the second ionisation energy (IE2), and write the equation of IE2 for calcium

A

The energy requires to remove one mole of electrons from one mole of gaseous 1+ ions to form one mole of gaseous 2+ ions, under standard conditions

24
Q

Define the third ionisation energy (IE3), and write the equation of IE3 for calcium

A

The energy required to remove one mole of electrons from one mole of gaseous 2+ ions to form one mole of 3+ ions, under standard conditions

25
Q

Why do the successive ionisation energies of an element increase

A
  • Once the first electron has been removed, a positive ion is formed
  • Removing another electron from a positive ion is more difficult that from a neutral atom
  • As more electrons are removed, the attractive forces between the electrons and nucleus increase due to the decreased shielding and increased proton to electron ratio
  • Increasing ionisation energy
26
Q

What factors affect the size of first ionisation energy

A
  • Size of the nuclear charge
  • Distance of outer electrons from the nucleus
  • Sheilding effect of inner electrons
  • Spin-pair repulsion
27
Q

Why does ionisation energy increase across a period

A
  • Nuclear charge increases along periods
  • This causes the atomic radius of atoms to decrease as the outer shell is pulled closer to the nucleus, so distance between valance electrons and nucleus decreases
  • Shielding remains constant as electrons are being added to the same shell
  • It become harder to remove an electron across a period as more energy is needed
28
Q

Why is there a decrease in IE1 between beryllium and boron

A
  • The fifth electron in boron is in the 2p sub shell, which is further away from the nucleus than the valance electron of beryllium which is in the 2s sub shell
  • Beryllium has a higher first ionisation energy due to the decreased atomic radius, meaning stronger attractive forces as decreased shielding
29
Q

Why is there a decrease in IE1 between nitrogen and oxygen

A
  • Due to spin pair repulsion in the 2p orbital of oxygen
  • Nitrogen has a higher first ionisation energy, as it has 1 electron in each 2p orbital
  • Oxygen has a lower first ionisation energy, as it has 2 electrons in one of the 2p orbitals, and the repulsion between the electrons makes it easier for them to be removed
30
Q

Why is there a decrease in ionisation energy between the last element in a period and the first element in the next period

A
  • There is increased distance between the nucleus and the valance electrons, as a new shell has been added
  • This new shell increases shielding
  • These two factors outweigh the increased nuclear charge
31
Q

What is the trend in ionisation energy down a group

A

Ionisation energy decreases down a group of

32
Q

Why does ionisation energy decrease down a group

A
  • The number of protons in the atom is increased, so nuclear charge increases
  • Atomic radius of atoms increase as more shells are added, causing distance between the nucleus and valance electrons to increase
  • This causes shielding to increase
  • These factors outweigh the increased nuclear charge, so it becomes easier to remove the valance electron
33
Q

How and why do electrons move energy level

A
  • Electrons move rapidly around the nuclei in energy shells
  • If their energy is increased, they can jump to a higher energy level, absorping energy
  • This can be reversed, and when they move down to their original energy levels they emit the energy they absorped
34
Q

What is a line emission spectrum, and why is there convergence

A
  • If the energy emitted by electrons moving down energy levels is visible, is can create a line emission spectra
  • Each line is a specific energy value, suggesting electrons can only posses a limited choice of energies (quanta)
  • Lines getting closer together in the blue end is called convergence, where lines converge to the higher energy, showing the electron reaching a maximum amount of energy
  • This maximum corresponds to the ionisation energy of the electron, n=2
35
Q

Describe what this graph shows

A
  • Large jumps show change in shells
  • Little jumps show change in sub shells
  • 1st electron is removed easily and has a low IE
  • 2nd electron is more diffficult to remove as there is no spin-pair repulsion
  • 3rd electron is harder still as it is in a different quantum shell (3p) so therefore closer to the nucleus
  • 4th electron is harder still as there is no spin-pair repulsion to help
  • Etc
36
Q

Define electron configuration

A

The arrangement of electrons in an atom

37
Q

What are orbitals

A

Atoms have sub shells to hold electrons, and within those shells are orbitals which hold the electron pairs

38
Q

What are principle quantum numbers

A
  • Used to number the energy levels or quantum shells
  • Low principal quantum numbers mean the shell is closer to the nucleus’s
  • Principle quantum numbers have fixed numbers of electrons it can hold, equal to 2n^2
39
Q

What are the number of orbitals in each sub shell

A
  • s: one orbital (holds 2 electrons)
  • p: three orbitals (holds 6 electrons)
  • d: five orbitals (holds 10 electrons)
  • f: seven orbitals (holds 14 electrons
40
Q

What is the s orbital shape

A
  • Spherical shape
  • Size of the s orbitals increases with increasing shell number
41
Q

What is the p orbital shape

A
  • Dumbbell shape
  • Orbitals occupy the x,y, and z axes and point at right angles to each other (perpendicular orientation)
  • Lobes of the orbitals become larger and longer with increasing shell number
42
Q

How do electrons fill orbitals

A
  • Electrons spin either clockwise or anti-clockwise, creating a N-S pole pointing up or down
  • Electrons with the same spin repel each other (spin-pair repulsion)
  • So electrons with occupy separate orbitals to minimise repulsion, and pair up with an electron with opposite spin
43
Q

What is Hund’s Rule

A

If there are only 3 electrons in the p shell, instead of two of them pairing up, they will all occupy separate orbitals for maximised stability

44
Q

What is the Pauli Exclusion Principle

A

An orbital can only hold two electrons, and they must have opposite spin, as the energy require to jump to a higher empty orbital is greater than the inter-electron repulsion, so electrons pair up in the lower energy levels first

45
Q

Why is the 4s orbital filled before the 3d orbital

A

As the 3d orbital has a higher energy than the 4s orbital

46
Q

Why are orbitals in the same sub shell degenerate

A

They have the same energy

47
Q

What is chromium’s exception to electronic configuration, and why

A
  • It is [Ar] 3d^5 4s^1 NOT [Ar] 3d^4 4s^2
  • It is more energetically stable
48
Q

What is coppers exception to electronic configuration, and why

A
  • It is [Ar] 3d^10 4s^1 NOT [Ar] 3d^9 4s^2
  • It is more energetically stable
49
Q

What is periodicity

A

The elements across periods which show repeating patterns in chemical and physical properties

50
Q

Why does the melting point in period 3 increase to silicon

A
  • Na, Mg, Al are metallic elements which form positive ions held together by a sea of delocalised electrons
  • Na donates its one valance electron to the ‘sea’, and Mg and Al donate their multiple valance electrons to the ‘sea’, causing the metallic bonding in Al to be stronger than Na
  • This is due to the electrostatic forces between a 3+ charge and the large number of e- is larger compared to 1+, causing Al to have a higher melting point than Na
  • Silicon has the highest melting point due to its giant molecular structure, where each Si atom is held to each other by strong covalent bonds
51
Q

Why does the melting point of period 3 decrease after silicon

A
  • P, S, Cl, and Ar are all non-metallic elements and exist with simple molecular structures held together by strong covalent bonds within molecules
  • Despite the strong covalent bonds within molecules, the molecules are held together by weak instantaneous dipole-induced dipole forces, which require little energy to break
  • Therefore the melting point decreases
52
Q

Why does atomic radius decrease across periods

A
  • Due to the number of protons (nuclear charge) and the number of electrons increasing by one each time you move an element to the right
  • Elements in a period all have the same number of shells (same shielding), so when you travel across a period, the nucleus attracts the electrons more strongly, decreasing the atomic radii
53
Q

Why is there a decrease in ionisation energy between magnesium and aluminium

A
  • The 13th electron in aluminium is in the 3p sub shell, which is further away from the nucleus than the 3s sub shell of magnesium
  • Magnesium has a higher IE, as the electrons are closer to the nucleus so require more energy to be removed
54
Q

Why is there a decrease in IE1 between phosphorus and sulphur

A
  • Due to spin-pair repulsion in the 3p orbital of sulphur
  • Phosphorus has higher ionisation energy as each of the electrons in its 3p sub shell are in separate orbitals, so there is no spin-pair repulsion to make the removal of an electron easy, like for sulphur

chemistry (19 decks)

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