CHAPTER 2: ELEMENTS + PERIODIC TABLE

Question 1

what is inside an atom?

Answer 1

• nucleus has positively charged protons and neutral neturons
• cloud containing negatively charged electrons form around nuclues
• electrons are bound to the nucleus by electrostatic attraction to the protons in the nucleus (opposite charges)

Question 2

properties of subatomic particles

Answer 2

• electrons are size 1/1800 relative to protons
• atoms are identified by how many protons they have
• an element is made up of atoms with the same number of protons

Question 3

atomic number

Answer 3

number of protons in the nucleus

Question 4

atomic mass

Answer 4

number of protons and neutrons in the nucleus

Question 5

isotopes

Answer 5

• changing number of neutrons
• are variants of a particular element which differ in neutron number
• all isotopes of an element have the same number of protons but diff number of neutrons
• have identical chemical properties: react the same way as other isotopes of a given element
• different physical properties
  
  • melting/boiling point
  • density
  • freezing point

Question 6

ions

Answer 6

• goal of every atom is to attain stability → neutral atoms w/ net overall charge of 0
• cations - positive ions
  
  • neutral atom loses an electrons
  • no. of electrons < no. of protons
  • net overall charge > 0
• anions - negative ions
  
  • neutral atom gains electrons
  • no. of electrons > no. of protons
  • net overall charge < 0

Question 7

molecules vs compounds

Answer 7

• molecules
  
  • group or cluster of 2 or more atoms held together by chemical bonds
  • atoms can be the same or different
• compounds
  
  • are substances made up of two or more diff type of atoms
• all compounds are molecules but not all molecules are compounds

Question 8

types of bonding

Answer 8

• metallic: metal + metal
• ionic: metal + non-metal
• covalent: non-metal + non-metal
  
  • all bonds involve electrons + all bonding involve changes to the number of electrons in the valence shell

Question 9

ionic bonding

Answer 9

• occur when there is a TRANSFER of electrons from a METAL to a NON-METAL
• metals tend to lose electrons and become cations
• non-metals tend to gain electrons and become anions
• eg. SODIUM CHLORIDE
  
  • sodium atom loses one electron → + ion
  • chlorine atom gains the electron from sodium → - ion
  • the 2 oppositely charged ions attract each other forming NaCl

Question 10

multiple valencies

Answer 10

• electrovalency = charge of ion
• some transition metals can form ions w/ diff valencies
• for compounds with these metals, roman numeral are used to specify charge
  
  • eg. iron (II) chloride

Question 11

polyatomic ions

Answer 11

• contains 2 or more atoms
  
  • fixed ratio
  • behave as a single unit w/ overall charge
  • subscripts are used to indicate internal ratio
• when more than 1 polyatomic ion is used in compound, brackets are required

Question 12

energy level in bohr

closest/furthest from nucleus

Answer 12

• electrons aren’t evenly spread but exists in layers called shells
  
  • shells can be called energy levels
  • closer to nucleus, the lower the energy level of electron
• electrons closest to the nucleus
  
  • highest force of electrostatic attraction to the nucleus
  • low kinetic energy - can’t move around as much
  • lowest energy level n=1
• electrons furthest from nucleus
  
  • lowest force of electrostatic attraction to the nucleus
  • high kinetic energy
  • highest energy level

Question 13

bohr model

electrons

Answer 13

• electrons revolve around nucleus in fixed circular orbits
• these orbits correspond to specific energy levels
• electrons can only occupy fixed energy levels → can’t exist between 2 energy levels
• orbits of large radii → higher energy levels

Question 14

emission spectra

Answer 14

• when atoms are heated, they emit light
• when light passes through a prism it produces a spectrum made of thin lines of different colours (emission spectra)
• each element has a unique emission spectrum

Question 15

how does bohr model predict emission spectra

Answer 15

• heating an element can cause an electron to absorb the energy and jump to a higher energy level
• shortly afterwards, the electron returns to its original level
• a fixed amount of energy is emitted as light (photon)
• ground state: lowest energy state of atom
• excited state: when electrons absorb energy + jump to higher energy level
• EVERY LINE IN THE EMISSION SPECTRA CORRESPONDS TO A SPECIFIC ELECTRON TRANSITION BETWEEN 2 SHELLS

Question 16

ionisation energy

Answer 16

• amount of energy required to remove an electron from an atom
• outer shell electrons: low ionisation energy → easiest to remove
• inner shell electrons: high ionisation energy → hard to remove

Question 17

electronic configuration of bohr/shell model

Answer 17

• electrons will fill shells nearest the nucleus first
• 1st shell holds 2
• 2nd shell holds 8
• 3rd shell holds 18
  
  • when filling 3rd shell, you must fill 8 first, then 2 in the 4th shell, then remainder on the third
• electron shell = n, max no. of electrons = 2n^2

Question 18

shortcomings of bohr model

Answer 18

1. couldn’t accurately predict emission spectra of atoms w/ more than one electron
2. unable to explain why electron shells can hold 2n^2 electrons
3. didn’t explain why the 4th shell has to accept 2 electrons first before filling the 3rd shell

Question 19

schrodinger model

Answer 19

• developed in 1926 by erwin schrodinger
• major energy level in an atom → shells
• shells contained separate energy levels of similar energies → subshells (s,p,d,f)
• each subshell is made up of smaller components → orbitals
  
  • total number of orbitals is n^2
  • each orbital contains 2 electrons

Question 20

difference between bohr and schrodinger model

Answer 20

• bohr: electrons travel along fixed paths
• schrodinger: electrons could be anywhere within the space

Question 21

electronic configuration of schrodinger model

Answer 21

• coefficient → shell number
• letters → occupied subshells
• superscript → number of electrons in the subshell

Question 22

condensed electronic configuration

noble gas notation

Answer 22

• noble gas notation
• symbol of noble gas is written in square brackets
  
  • eg. [Ne]
• this emphasises valence shell electrons
• example: phosphorus
  
  • [Ne] 3s^23p^3

Question 23

exceptions to schrodinger model

Answer 23

CHROMIUM (Cr) 24

should be: 1s2 2s2 2p6 3s2 3p6 3d4 4s2

actual: 1s2 2s2 2p6 3s2 3p6 3d5 4s1

• the 3d5 4s1 is slightly more stable as each of the 5 3d orbitals are exactly half-filled

COPPER (Cu) 29

should be: 1s2 2s2 2p6 3s2 3p6 3d9 4s2

actual: 1s2 2s2 2p6 3s2 3p6 3d10 4s1

• 3d10 4s1 has 5 completely filled orbits rather than 3d9 4s2

Question 24

modern periodic table

Answer 24

arranged in order of increasing atomic number

vertical columns → groups 1-18

horizontal rows → periods 1-7

main group elements are in groups 1,2, 13 -18

transition metals: group 3-12

Question 25

groups

Answer 25

• for main group elements, group number determines the number of valence electrons
  - same number of valnce electrons = similar properties
  
  • group 1: alkali metals → 1 valence electron (highly reactive)
  • group 17: halogens → 7 valence electrons (highly reactive)
  • group 18: noble gases → 8 valence electrons (least reactive elements)
    
    • exception: helium is a noble gas

Question 26

periods

Answer 26

tell us how many occupied shells an atom has

Question 27

blocks

Answer 27

4 main blocks

• elements in each block have the same type of subshell (s,p,d,f) as their highest energy shell
• so elements in s-block have s-subshell has highest subshell

Question 28

critical elements

Answer 28

• elements that are heavily relied on for industry and society
  - eg. electronics, food supply
  
  • over 40 elements are identifies as endangered elements
    
    • small deposits rapidly disappearing
      
      • iridium, platinum, osmium
    • little to no recycling + recovering of the elements
    • supply centred in war + conflict areas
      
      • gold, tin, tungsten
  • requires new ways to recover and recycle endangered elements

Question 29

core charge (effective nuclear charge)

Answer 29

• is a measure of the attractive forces felt by the valence electrons towards the nucleus
  
  • in atoms with 2 or more shells, the attraction between valence shell and nucleus is reduced by repulsion between inner shell electrons and the valence electrons
• core charge = number of protons in the nucleus - number of total inner-shell electrons

Question 30

core charge trend

Answer 30

down a group:

• core charge remains constant
• valence electrons are held less strongly as they are further apart

left to right across a period:

• core charge increases
• valence electrons are more attracted to the nucleus as core charge increases

Question 31

electronegativity

Answer 31

• ability of an atom to attract electrons TOWARDS ITSELF
• positive pull comes from nucleus, so if core charge increases, electronegativity increases
  
  • electrons want to be as close to nucleus as posible

Question 32

electronegativity trend

Answer 32

down a group:

• trend in electronegativity decreases
  
  • v.e are less attracted to the nucleus as they are further from the nucleus (more shells down a group)

across a period:

• trend in electronegativity increases
  
  • v.e are more strongly attracted to the nucleus as core charge increases

Question 33

atomic radius

Answer 33

• the distance from the nucleus to the valence shell elecrons
  
  • atoms with high electronegativity or core charge pull electrons closer to the nucleus (becoming smaller atoms)

Question 34

atomic radius trend

Answer 34

down a group

• atomic radii increases
  
  • the number of shells increases

across a period

• decreases
  
  • core charge increases, electrons are pulled closer to nucleus → smaller atomic radii

Question 35

first ionisation energy

Answer 35

• the energy required to remove one electron from an element in the gas phase
  
  • magnitude of ionisation energy reflects how strongly the valence electrons are attracted to the nucleus
  • high f.i.e means electrons are harder to remove

Question 36

first ionisation energy trend

Answer 36

down a group

• first ionisation energy decreases
  
  • valence electrons are less attracted to the nucleus as they are further from the nucleus

across a period

• first ionisation energy increases
  
  • the core charge increases, so v.e are more attracted to nucleus and harder to remove

Question 37

reactivity

Answer 37

an indication of how easily an atom of that element loses or gains electrons

Question 38

metal reactivity

Answer 38

• moving down group, no. of shells increases
• electrostatic force of attraction between v.e and nucleus become weaker
• reactivity of metals increases down a group, but decreases across a period

Question 39

non-metal reactivity

Answer 39

• the more easily a non-metal gains an electron, the more reactive it becomes
• the more electronegative an element, the more readily it will attract electrons
• reactivity of non-metals increases across the period but decreases down the group

Question 40

metallic character

Answer 40

• metals conduct electricity and are usually solid at room temp
• non-metals don’t conduct electricity and are usually gases at room temperature
• metalloids are located between metals and non-metals
  
  • exhibit both metallic and non-metallic properties