Flashcards
3.1 The Periodic Table
Cards in this deck (21)
- 1Describe the trends across a period - ionisation energy values show a general increaseAcross Each Period:
- number of protons in the nucleus increase = higher attraction on the electrons
- electrons are added to the same shell = outer shell is drawn in closer
- same number of inner shells = electron shielding will hardly change - 2Describe the trends down a group - 1st ionisation energy decreases moving down the group- the number of shells increase so the distance of the outer shell from the nucleus increases
- more inner shells = shielding effect on the outer electrons from the nuclear charge increases = weaker attraction - 3Describe the melting and boiling points of giant metallic structuresHIGH M+B Points
- electrons are free to move throughout the structure, but positive ions remain where they are
- attraction between positive ions and negative electrons is very strong
= high temperature needed to overcome the metallic bonds and dislodge the ions from their rigid positions in the lattice - 4What is ductility in metals? What does malleable mean?Ductile = can be drawn out or stretched
Malleable = hammered into different shapes - 5Describe the trends in melting points between group 1 and 14increase steadily - elements have giant structures, nuclear charge increases with number of electrons = stronger attraction
- 6Describe the trend in melting points for group 14 to 15sharp decrease in melting point
elements have simple molecular structures - each individual molecule is attracted to others by weak inter molecular forces - 7Describe the trend in melting points for 15 and 18melting points remain relatively low
- elements have simple molecular structures - 8Describe some properties of the Halogens (Group 7)- low melting and boiling points
- exist as diatomic molecules, X^2 - X = the halogen. Eg, Br^2 - 9Describe the trend in boiling points for the Halogens (Group 7)moving down the group = increased boiling point
- each successive element has an extra shell of electrons
= higher level of London forces ( induced dipole-dipole forces) between the molecules - 10Reactivity of the Halogensvery reactive and electronegative
- good at attracting and capturing electrons - good oxidising agents. Electrons are gained = 1- ions (halide ions)
-reactivity and oxidising power decreases down the group because:
- atomic radius increases (nuclear pull further away from incoming electrons)
- electron shielding increases
- ability to gain an electron in the p sub shell and form 1- ions decreases - 11Explain the successive ionization energy theoryEach successive energy is higher than the 1 before
- as each electron is removed, less repulsion between the remaining electrons = each shell is drawn in closer to the nucleus = stronger nuclear charge
- positive nuclear charge will outweigh the negative charge each time an electron is removed
- as the distance between each electron and the nucleus decreases (atomic radius), nuclear attraction increases - more energy needed to remove each successive electron - 12How does electron shielding affect ionization energyinner shells of electrons repel the outer shell electrons because all negative
- repelling effect = electron shielding
- more inner shells = more electron shielding, smaller nuclear attraction experienced by outer electrons - 13Explain how nuclear charge affects ionization energy?the higher nuclear charge = larger the attractive force on outer electrons
- 14Explain how atomic radius affects ionization energy?the larger the atomic radius = smaller nuclear attraction
= positive charge of the nucleus is further away from the outermost electrons - 15What are the factors that affect ionization energy?nuclear attraction experienced by an electron depends on:
- atomic radius
- nuclear charge
- electron shielding/screening - 16Why are metals good electrical conductors?the delocalized electrons can move freely throughout the metallic lattice structure
= allows metals to conduct electricity, even in the solid state - 17Describe the ionization energy trends between groups 2 to 13DECREASES between 2 to 13
- group 13 elements have outermost electrons in p orbital,
whereas group 2 have outermost electrons in an s-orbital
- p-orbitals have higher energy than s so are further away from nucleus = slightly easier to remove = lower ionization energies - 18Describe the ionization energy trends between group 15 to 16DECREASES between 15 to 16
- from 13 to 16, outer electrons are found in p-orbitals
- 13,14,15 - p orbitals contain only ONE electron whereas in 16, the outermost electrons are spin paired in the p - more repulsion
= outer electron easier to remove = lower ionization energy - 19Describe the general ionization trend across a periodsharp decrease between ionization energy between end of 1 period + start of the next
- reflects additions of new shell - further from the nucleus
- increase in atomic radius + electron shielding - 20Describe the ionization energy trends down a groupDECREASE
- number of shells increase, so atomic radius (distance between outer electrons from nucleus) increases
= weaker force of attraction (electrons are not as close to the nucleus)
- more inner shells = greater electron shielding, from the outer electrons + nuclear charge
= weaker force of attraction
= less energy needed to remove electron
- number of protons also increase but resulted increased attraction from the increased atomic radius and shielding is outweighed - 21—
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- A level Chemistry (Module 3 - Periodic Table + Energy) Flashcards on 3.1 The Periodic Table, created by Amna Jamil on 22/05/2017.
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- Amna Jamil
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