CHEMISTRY FLASHCARDS

UNIT 2 - 56

Does the energy increase or decrease as the electrons get further from nucleus?

• energy increases

What are the 7 diatomic molecules?

1. H_² (Hydrogen)

2. N₂ (Nitrogen)

3. O₂ (Oxygen)

4. F₂ (Fluorine)

5. Cl₂ (Chlorine)

6. Br₂ (Bromine)

7. I₂ (Iodine)

Do all electrons in a atom have the same energy? Why?

• They have different energies as the further from nucleus the more energy

What are Isotopes?

• Atoms with same number of protons

• different neutrons

What does the symbol A indicate? How do you calculate it?

• Atomic number

• Protons + neutrons

What are the three subatomic particles?

• protons

• neutrons

• electrons

What does the symbol Z indicate?

• The number of protons

Whats the difference between negative ions and positive ions?

• Negative - More electrons than protons (- = gain) - compared to neutral atoms

• Positive - Less electrons than protons (+ = lost) - compared to neutral atoms

Whats the difference between a period number and group number on the periodic table?

• period - row (number of energy levels (shells))

• group - column group - number of valence electrons

how many electrons does the 1st, 2nd and 3rd shell contain? What model is this called?

• 1st = 2

• 2nd = 8

• 3rd = 18

• Bohr’s Atomic Model

fill in the table:

Shell	Subshell	Maximum electrons	Maximum subshell	Maximum orbital

Shell	Subshell	Maximum electrons	Maximum subshell	Maximum orbital
1	1s	2	2	1
2	2s, 2p	8	6	3
3	3s, 3p, 3d	18	10	5
4	4s, 4p, 4d, 4f	32	14	7

What is the electron configuration order and its energy order - sub-shells number of orbitals and electrons?

Which shells has the highest and lowest energy?

• s (1 orbital, max 2 electrons) - lowest energy

• p (3 orbitals , max 6 electrons)

• d (5 orbitals , max 10 electrons)

• f (7 orbitals , max 14 electrons) - highest energy

Why does the orbital energy increase?

• decrease in effective nuclear attraction - reduces the nucleus’s control to hold the electron

• shielding and distance increases - increases energy

What are the 3 principals? provide an example for each

1. Aufbau principal - electrons must fill lower atomic and energy orbitals before occupying the higher ones. (front of lecture theatre to back of theatre)

2. Pauli exclusion principal - hold max of 2 electrons with opposite spins. Seesaw two people sit but one goes spins/goes up and the other spins opposite/ goes down.

3. Hund’s rule - electrons occupy first orbital at a time before pairing. everyone gets own work, before pairing up

What is ionisation energy?

• energy required to remove one valence electron

Why does ionisation energy increase across the same shell (period)?

• across period - nuclear charge increases causing the electrons to be more attracted to the nucleus making it harder to remove. Whilst, the shielding remains relatively similar.

Why is there a larger jump after valence electrons are removed?

• Next electron is in the inner shell and is closer and more attracted to the nucleus

How does successive ionisation energy indicate valence electrons? for example, what does it indicate when its after the 3rd electron?

• Number of small increases before a dramatic increase = number of valence electrons

• That the atom has 3 valence electrons

Why do Isotopes of same element behave chemically similar? But not physically?

Chemically:

• Same number of electrons and similar bonding and electron configuration

Physically:

• due to different neutron numbers it changes its physical properties, such as density and mass

What does “X” stand for in Isotopes? (^{A_Z X)}

• X = Element symbol

What does “A” stand for in Isotopes?

• Mass number

Why is the weighted average often recorded for an isotope?

• Takes into account of weight and abundance of each isotope

What does abundance mean? provide an example

• Is the amount of something that exists compared to the rest.

• For example C-35 (75%) and C-37 (25%). C-35 is more abundant as it appears more often in nature.

Why is RAM not a whole number? Why do some elements have fractional RAM values?

• Weighted average of different isotopes contain different mass and abundance

• Many elements exist as a mixture of different Isotopes - so average is not whole.

Group metals and gases Table - Perodic trends and tables. Fill in the table below

Group:	Name:	Properties:	Valence electrons
		-
		-
		-
		-

 

Group:	Name:	Properties:	Valence electrons
1	Alkali	-     One valence electron, Soft metals, Most reactive and reacts with water	From group 1 to 18, it decreases in reactivity
2	Alkali earth metals	-             2 Valence electrons, Hard metals, Reacts slower than group 1
17	Halogens	-    7 valence electrons, non-metal, Diatomic molecules, and Reacts highly with 1 and 2 metals. -              complete outer shell, little to no in reaction.
18	Nobel gases

 

What is the shielding effect?

• Inner electrons block the nucleus’s ability to pull on the outer electrons

What is electronegativity? and does an increases/ decrease in electronegativity indicate a high/low attraction? answer for both an increase and decrease in energy

• Ability of atom to attract shared electrons while being bonded to another atom

• increase in electronegativity = high attraction

• decrease in electronegativity = low attraction

What are Transitional elements?

• Elements that form ions due to various oxidation states (loss or gain of electrons)

• found in d-block

• forms one or more ions - as it loses different numbers of valence electrons

Properties of Metal and non-metal include? Also why metals conduct electricity

Metal:

• Good conductors of heat and electricity - contains delocalised electrons which move around and carry charge.

• Ductile (shaped without breakage)

• Malleable

• Lustrous (Shiny)

Non-Metals:

• Poor conductors of heat and electricity - no delocalised or free electrons - doesn’t carry charge

• dull (not shiny)

• brittle - breakable

What are the properties for metallic? Types of bonding (1/4)

• high melting and boiling points

• strong bond between delocalised electrons and metal ions

• conducts electricity in liquid and solid

What are the properties for ionic? Types of bonding (2/4)

• No conduction of electricity in solids but does in liquids

• Ions are fixed in solid but free in liquid

• high melting and boiling points

• strong electrostatic attraction between opposite charged ions

What are the properties for covalent molecular? Types of bonding (3/4)

• low melting and boiling point

• weak intermolecular force

• doesn’t conduct electricity in liquid or solids

• no free or delocalised ions

What are the properties for Giant Covalent network? Types of bonding (4/4)

• extremely high melting and boiling point

• does not conduct electricity

• no charged particles

What are Ions?

Ions are a group of atoms that have lost or gained electrons

• loses electron - positive ion

• gains electron - negative ion

What is electrostatic attraction?

• Type of force that pulls and holds oppositely charged particles together

What are the oxide changes and why across period 3? Why?excluding the noble gas

• Basic → Basic (if it produces OH^-) → Amphoteric→ Rest is Acid (if it produces H⁺)

• this is due to metals (Basic) easily losing electrons and forming Ionic compounds.

• whereas, non-metals (acidic) share electrons and form covalent oxides.

What are compounds?

• products formed in a reaction

What particle attraction does Ionic bonding attract?

• oppositely charged particles

What is covalent bonding?

• sharing of electrons

What does valency mean? and what does the number of valency indicate?

• Number of bonds an atom can form

• number of valency = number of covalent bonds

molecular prefixes

Number of atoms:	Prefix-	Number of atoms:	Prefix-
1		6
2		7
3		8
4		9
5		10

Number of atoms:	Prefix-	Number of atoms:	Prefix-
1	Mono-	6	Hexa-
2	di-	7	Hepta-
3	Tri-	8	Octa-
4	Tetra-	9	Nona-
5	Penta-	10	Deca-

What is are the types of formulas? (Empirical and molecular?)

• Empirical formula - simplified whole number ratio of atoms in a compound

• molecular formula - shows the actual number of each atom in a molecule

Whats the difference between a lone and a bonding pair (lewis dot structure)?

• lone pair = no sharing of electrons (dots (..))

• Bonding pairs = sharing of electrons (lines (-))

What is a pure substance and what are two types?

• Elements and compounds

• pure substance is one type of particle

What’s the difference between Homogeneous and Heterogeneous mixtures?

• Homogeneous mixture - particles evenly mixed and throughout visibly same

• Heterogeneous mixture - particles are unevenly mixed and parts are visibly different

What does Lattice mean?

• Fixed 3D repeated pattern of a particle in a solid.

What is an Allotrope? and What is a phase change? provide an example of each.

• Allotrope - same element and physical state but different structural form - carbon: diamond and graphite

• Phase change - same substance changes in different physical state (solid, liquid and gas)

What is the difference between exothermic and endothermic reactions? provide an example for each. Also are the products energy higher or lower than its reactants?

Exothermic reaction:

• Is the exist of heat from the system into its surroundings. E.g. Campfire.

• Causes the net energy to be released so the products heat is lower than reactants

• Negative heat and feels warm

Endothermic reaction:

• Is the absorption of heat from the surroundings into the system. E.g. heat packs

• Energy is absorbed so products heat is higher than reactants

• Positive heat and feels cold

What happens to the energy when bonds break and bonds form?

• Bonds break - absorption of energy

• Bond forming - release of energy

What does activation energy and (triangle/delta symbol) H indicate?

• Activation energy is the minimum amount of energy required to start/produce a reaction or breakage of bonds

• (triangle/delta symbol) H indicates the enthalpy change (energy change) during a reaction.

What is a bond enthalpy? And why is it measured in averages?

• activation energy to break one mole of a bond in a gas state

• Average values are taken from different molecules

What is a mole? Refer to Avogadro’s number. Also why is moles useful

• fixed substance containing 6.02 × 10²³ particles

• Count particles by measuring its mass

U2

What are the five molecular shapes? What are its characteristics: degree, bonding electrons, lone pairs?

• Linear: 180°, 2 bondings and 0 lone

• Bent: 104.5°, 2 bonding and 1 lone

• Trigonal Planar: 120°C, 3 bonding and 0 lone

• Trigonal Pyramidal: 107°, 3 bonding and 1 lone

• Tetrahedral: 109.5°C, 4 bonding and 0 lone

Central atom	Typical bond	Lone pair	Shape	Angle
Group 2
Group 13
Group 14
Group 15
Group 16

Central atom	Typical bond	Lone pair	Shape	Angle
Group 2	2	0	Linear	180°
Group 13	3	0	Trigonal planar	120°
Group 14	4	0	Tetrahedral	109.5°
Group 15	3	1	Trigonal pyramidal	107°
Group 16	2	2	Bent	104.5°

What does an increase and decrease in lone pairs result in? also for the central atom?

• increase = small angles

• More lone pairs on the central atom produces a greater repulsion (due to occupying more space) against bonding pairs - decreasing angles

What do polar molecules contain?

What determines an atoms pull/strength when polar?

• Electronegativity = strength of the atoms ability to attract bonding electrons

• Example NH_³ N has a higher electronegativity than H = uneven electron sharing = uneven charge distribution

Compare polar and non-polar molecules - discuss dissolving (What it can be dissolved in)

• Both influenced by an atoms Molecular structure

• Can only dissolve by another molecule in the same classification. For example polar can only dissolve in polar.

• Differences:

Polar	Non-polar
-              Lone pairs -              Not symmetrical	-              No lone pairs -              Symmetry

Compare Intramolecular and Intermolecular force

• Both are forces and create an attraction between the molecules

• Intramolecular - force inside molecule

• Intermolecular - for between molecules

What are the three intermolecular forces (lowest strength to highest strength)? What does each do?

• Dispersion forces - Occurs in all molecules. Caused by an uneven temporary distribution

• Dipole-Dipole forces - Occurs in only polar molecules and creates a permanent uneven distribution

• Hydrogen bonding - Occurs when FON molecules bond with hydrogen

Strong vs Weak IMF:

Strong IMF:	Weak IMF:
-              High melting and boiling point -              Low vapour pressure (ability for molecule to escape in a gas phase)	-              low melting and boiling point -              High vapour pressure (ability for molecule to escape in a gas phase)

How does the dispersion force increase? How does the boiling force increase?

• More electrons → more distance → bigger dispersion force

• Stronger attractive force → increase in boiling points

what does symmetrical and non-symmetrical mean?

• Symmetrical - Uneven distribution of charges

• Non-symmetrical - Even distribution of charges

What forces are polar or non-polar?

Non-polar - dispersion

Polar - Dipole- Dipole and Hydrogen

How does a intermolecular forces affect melting point, vapour pressure - why? What does a boiling point suggest about the vapour pressure

• high melting

• low vapour - fewer molecular escape into gas-phase

• it is low

What is the strength of dispersion forces influenced by?

• Molecular Mass

• Chain length of molecule

• Molecular shape

What is vapour pressure in both forms of high and low?

• Substance’s liquid particles trying to escape to in the form of a gas-state

• high - large amount of liquid particles trying to escape in the form of a gas

• low - fewer to no liquid particles trying to escape in the form of gas particles

What does the “like dissolve like” mean in polar and non-polar

• Polar dissolves in polar solvents

• Non-polar dissolve in non-polar solvents

What is a capillary action?

Ability for a liquid to oppose gravitiy to allow it to flow through narrow spaces

Why is water a polar molecule?

O has a higher electronegativity than the 2 hydrogens. due to the uneven distribution of charges = polar

What is a solvent, solute, and solution?

• Solvent - substance that dissolves (does)

• Solute - substance being dissolved in a solute (gets)

• Solution - Homogenous mixture

Difference between strength, concentration, and dilute?

• Strength - describes how much a substance is acidic or basic

• concentration - describes the amount of solute per vol of a solution

• Dilute

What do square brackets suggest?

[ ] concentration of species in mol L ^-1

difference between concentrate and dilute solutions

• contains more particles per volume of solution

• contains fewer particles per volume of solution

difference between unsaturated, saturated, and supersaturated? what happens to each? and positioning on line? Provide an example

Unsaturated - Holds max amount of solute possible. will instead sink to bottom and not dissolve. Below

Saturated - Holds a smaller amount of solute, than the maximum can hold. easily. On line dissolves

Supersaturated - Contains more solute than theoretically possible. Above line

E.g. Sugar

What is molarity concentration? Formula

• Amount of solute in moles per litre of solution

• Formula - c = n/V