Chem IB Exam Structure 1: Models of Particulate Nature of Substances

Define Chemical Element

It is Pure Substance made up of only ONE type of atom

Define atom

Smallest particle of an element to show the characteristic properties of that element. Example: Au (gold)

Define Compound

Chemical combinations of different elements

• They contain FIXED proportions/ratios of elements, held together by chemical bonds

• Ex: NaCl

Define mixture

Composed of Two or MORE substances in which no chemical combination had occurred

• It does not have a fixed composition.

Define Homogenous Mixture

Uniform composition and properties. It has the same properties even if sample is taken from anywhere. Ex: saltwater

Define Heterogenous Mixtures

Non-uniform composition and properties are NOT same throughout. Ex: salad

List methods of separating compounds

Filtration: Solid particles in a liquid are removed through the use of a filter that allows only fluid to pass

Recrystallization: A substance is separated by dissolving the mixture in a solvent and then crystallizing desired substance from solution

Distillation: Components of a liquid mixture are separated through selective boiling and condensation

Paper Chromatography: Separates components of a mixture based on absorption to a solid phase.

Understand

Define Temperature

Measure of average kinetic energy of particles of a substance

Fluids

Liquids and gases are fluids which refers to their ability to flow

Define sublimation and deposition

Sublimation: Direct inter-conversion from solid to gas without the liquid phase

Deposition: Change from gas directly to solid. e.g. frosting

Distinguish between Vaporization and Evaporation

Vaporization: State change of matter from liquid to gas and can be accomplished through evaporation or boiling at FIXED temperature

Evaporation: State change that occurs only at surface of a liquid and at a temperature below the boiling point eg. clothes drying outside

Boiling Point

When all molecules in a liquid have enough kinetic energy to change into a gas

What occurs in heating curve from a-b

The solid is heated and as vibrational energy of particles increases so does the temperature

 Note, during a state change there will be no increase or decrease in temperature

• Adding temperature only increases the kinetic energy of the molecules, which will eventually break the bonds, then the molecules will change state

What occurs in heating curve from b-c

It is the melting point. Vibrations are energetic enough for particles to move away from fixed positions and form a liquid Temperature remains constant at this point

What occurs in heating curve from c-d

Liquid is heated, particles gain energy and temp increases

What occurs in heating curve from d-e3

d-e: boiling point, sufficient energy to break all interparticle forces and form a gas.

• Requires more energy than melting

• Temperature remains constant

• Bubbles of gas visible throughout

volume of liquid

What occurs in heating curve from e-f

gas is heated under pressure, kinetic energy of particles continues to rise, temperature also rises

Conversion of Kelvin to Celsius

Temperature (K) = Temperature (°C) + 273.15

Distinguish between atomic # and Mass #

Atomic #: Number of protons in an atom

Mass #: Number of protons + the number of neutrons in an atom

Ions?

When atoms gain or lose electrons

Cations vs. anions

Cations are atoms that are positively-charged ion because they lose electrons

Anions are atoms that are negatively-charged ion because they gain electrons

Define isotopes

They are atoms of an element that have the same # of protons but different # of neutrons

• Have very similar chemical characteristics because # of protons and electrons are the same

• Have different atomic mass because of different # of neutrons which leads to different physical properties

Atomic radius vs. Ionic radius

Atomic Radius: The distance of the nucleus from its outermost electron

Ionic Radius: The radius of the atom that has gained or lost electron(s) and become negatively or positively charged

Relative Average Mass Formula

Define Mass Spectra and be able to interpret it

Mass spectrometer is an instrument used to measure mass of individual atoms.

• Results are presented as a mass spectrum, where % abundance is plotted against mass/charge ratio of different ions

• Relative average Mass can be calculated from this data

Define Electromagnetic radiation

A from of energy that propagates through space at the speed of light as electromagnetic waves AKA photons

Define Frequency

Number of waves that pass a point in 1 second

Define Emission Spectra

produced by atoms emitting photons when electrons in excited states return to lower energy levels.

Explain Convergence of lines on the hydrogen spectrum

As energy levels increases, they get closer together and converge at high energy.

Distinguish between a continuous and a line spectrum.

When excited electrons ‘fall’ from a higher to a lower energy state, photons with a discrete amount of energy are emitted. The emission spectrum of atoms is a line spectrum: only light of a particular colour (discrete energy) is emitted.

Describe the emission spectrum of the hydrogen atom, including the relationships between the lines and energy transitions to the first, second and third energy levels.

• Electrons that ‘fall’ to the groundstate (n = 1) emit photons with the greatest amount of energy (UV radiation). The length of the arrows is proportional to the amount of energy. Electrons that ‘fall’ to n = 2 emit visible light and to n = 3 emit infrared radiation.

line spectra converge at higher energy because the energy levels inside the atom are CLOSER together at higher energy.

• When an electron reaches the highest energy (n=infinity), the electron leaves the atom and results in an ion

Energy level

It is how far an electron is away from the nucleus and denoted as ‘n”

• The maximum number of electrons that can be found in each

  energy level can be found using 2n²

Sub-level

Shape of region in which electron can be found. Sublevels contain fixed # of orbitals

Orbitals

Region of space where there is a High probability of finding an electron

o    s: 1 orbital, 2 electrons

o    p: 3 orbitals, 6 electrons

o    d: 5 orbitals: 10 electrons

o    f: 7 orbitals, 14 electrons

What is the s sub-level

• Every s subshell consist of one spherical orbital, which is further away from the nucleus the higher the shell number.

What is a Node

A region with ZERO probability of finding an electron

What is the p sub-level

• The second energy level, n=2, can have s AND p sub-levels

• The p subshell always contains 3 orbitals which are aligned along the x,y, z-axis. Since the second shell (n = 2) contains the 2s 2p subshells which can host 2 and 6 electrons respectively, the maximum total number of electrons in the second shell is 8.

Apply the Aufbau principle, Hund’s rule and the Pauli exclusion principle to deduce electron

configurations for atoms and ions up to Z = 36.

• Electrons fill into the lowest energy orbitals first. Each orbital can hold at most 2 electrons (Aufbau principle)

• If two electrons occupy the same orbital, they must have a different spin (Pauli Exclusion Principle)

• If multiple orbitals have the same energy, they are filled singly first, then doubly. (Hund’s Rule)

Writing electron confi/arrangement. Also what are lectron config for Chromium and Copper

• The 4s orbital is filled first before 3d, but is removed first before 3

• The electronic configurations of the transition elements copper and chromium do not follow the expected patterns

  • Chromium has the electron configuration: [Ar] 3d5 4s1

  • Copper has the electron configuration: [Ar] 3d10 4s1

What is the ground state

When an electron is with no addition of energy

What is the infinity level

The highest possible energy level that an electron can have and still be part of the H atom

Define First Ionization Energy

The MINIMUM amount of energy required to remove 1 mole of electrons from 1 mol of gaseous atoms.

• X9G) → X⁺ + e^-

What factors influence 1st ionization energy

1. Size of the nuclear charge

2. Distance of outer electrons from the nucleus

3. Shielding effect

How does the size of nuclear charge affect 1st ionization energy

• As the # of proton increases, the nuclear charge increase

• The larger the positive charge, the GREATER the attractive electrostatic force between the nucleus and electrons (More protons create a stronger electrostatic force, pulling electrons closer to the nucleus.)

• SO, a larger amount of energy is needed to overcome these forces and remove an electron

• As the proton number increases, ionization energy increases: First ionization energy increases across each period

How does the distance of outer electron affect 1st ionization energy

• The electrostatic attraction decrease as distance increases so electrons in shells are more weakly attracted to the nucleus

• The further the outer electron is form the nucleus, the LOWER the ionization energy is

• Thus, ionization energies tend to decrease down a group of the periodic table

How does shielding effect affect 1st ionization energy

• It occurs when the full inner shells of electrons PREVENT electrons in higher levels from the pull of the nucleus

• The GREATER the shielding effect, the lower the electrostatic force between outer electrons and the nucleus

• Thus, ionization energy is lower as the number of full electron shells between the outer electrons and the nucleus increases

Formula for effective nuclear charge

Z_eff = Z - S

Z: atomic number(#of electrons/[protons)

S: shielding electrons