IB CHEM- STRUCTURE 1

Pure substances

Made of only one type of substance

Have a fixed composition

Mixtures

Combination of 2+ pure substances

NO fixed composition

• 2+ different elements or compounds NOT chemically combined

• Can be separated by physical methods (separation techniques)

• Keep the properties of the individual elements of compounds

• Alloys are included even though they contain metallic bonds

Homogenous [mixture]

• no visible phases or boundaries

• Ex. Saltwater

• uniform: different parts are equally distributed and in the same state

Heterogeneous

• visible phases or boundaries

• Ex. Oil and water

• different parts of the mixture have different compositions or states

Elements

• simplest form of matter

• One type of atom

• Cannot be chemically broken down into simpler substances

• Ex. Iron (Fe)

Compounds

• 2+ different elements chemically combined in a fixed ratio

• Cannot be separated by physical methods

• Ex. Methane, CH₄

• properties of compounds are much different from those of individual elements

Solute

A substance (usually solid) dissolved in a solution

Solvent

A substance (usually liquid) in which other substances are dissolved

Solution

A homogenous mixture composed of a solute dissolved in water (the solvent)

Dissolve

When solute particles are surrounded by solvent particles

Solubility

The ability of a substance to dissolve into a solvent to form a solution

Insoluble

When a solute does not dissolve in a solvent

Soluble

Can dissolve in a solvent to produce a solution

Filtrate

A substance that has passed through a filter

Residue

The insoluble component (usually a solid) of a mixture that remains on the filter paper after filtration

Volatility

The tendency of a substance to undergo evaporation

Distillate

The part of a liquid mixture that evaporates and condenses in the distillation process

Filtration

Physical property: particle size

• separates an insoluble solid from a liquid

• Heterogenous mixture

• Ex. Salt and vinegar

• Large particles get stuck on filter paper, small pass through

Evaporation

Physical property: boiling point

• separates a dissolved solid from a liquid

• Homogeneous mixture

• Ex. Salt and water

• Liquid evaporates leaving the solid behind

Distillation

Physical Property: boiling point

• separates two liquids

• Homogeneous mixtures

• Ex. Water and ethanol

• One liquid evaporates first, then condenses; pours into a separate container

Recrystallization

Physical property: solubility at different temperatures

• separates impurities from a solid

• Ex. Purify sugar crystals

• Impure mixture is dissolved in hot liquid; as it cools, pure crystals form and impurities stay dissolved; use filtration to separate crystals from impurities

• Least-soluble solution will crystallize first

Chromatography

another type of separation technique used to separate mixtures that contain very small amounts of each component, or to determine how pure a substance is

• set ups contain 2 phases

  • Mobile phase- moves 

  • Stationary phase- stays in place

• Works due to components of a mixture having different tendencies to either absorb onion a stationary surface or dissolve into a mobile solvent 

• 2 types: paper chromatography and thin-layer chromatography (TLC)

paper chromatography 

• a mixture is applied to a piece of chromatography paper

  • Paper is in stationary phase

  • Solvent is in mobile phase

• As the mobile phase starts to climb up the paper, the mixture will be carried with it each at a different rate

  • Components that have a great affinity for the mobile phase will move father up the paper because they will interact with the mobile phase more

  • Components that have a great affinity for the stationary phase will move less far up the paper because they will interact with the stationary phase more

Thin-Layer Chromatography (TLC)

• very similar to paper chromatography 

• Advantage: separated components can be recovered pure

Kinetic Molecular Theory: States of Matter 

Matter is composed of particles (if an element) or molecules (if a compound). These particles have kinetic energy (motion energy) so they are constantly moving 

• Higher temperature= more movement; more likely to be a gas (straight line motion)

• Lower temperature=less movement; more likely to be a solid (vibrational motion)

• Collision between particles are elastic (no loss in kinetic energy)

Solid-s

• fixed volume 

• Fixed shape 

• Cannot be compressed

• Strong attractive forces between particles 

• Particles vibrate in fixed positions it do not move around 

• More dense 

• Have least amount of kinetic energy

Liquid- l

• fixed volume

• No fixed shape

• Cannot be compressed

• Weaker attractive forces between particles

• Particles vibrate, rotate, and move around

• Less dense

• WILL TAKE THE SHAPE OF THE BOTTOM OF THE CONTAINER BUT NOT OF THE CONTAINER ITSELF

Gas-g

• no fixed volume

• No fixed shape

• Can be compressed

• Negligible attractive forces between particles

• Particles vibrate, rotate and move fast

• Least dense

• Have the most amount of kinetic energy 

Aqueous- aq

• a solid is dissolved in H₂O

• Each molecule or ion is surrounded by water

Solvation

The separation of a heterogeneous mixture of two solids based on difference in solubility if one of the substances is soluble in a solvent, but the other solid is insoluble. 

• during solvation, the solvent molecules (most  often water) surround the soluble molecules and dissolve the solid into a solution

• Th insoluble solution can now be separated by filtration

• The soluble substance can be separated from the solution by evaporation

Plasma

An ionized gas mainly found in outer space

Density 

Mass per unit volume 

• substances with higher densities will feel ‘heavier’ compared to substances with lower densities (of the same volume)

• Formula: d=m/v

• Solids ten times have higher densities than liquids which tend to have higher densities than gases 

Changes of state or phase changes 

Occur when a substance changes from one physical  state to another 

• a physical change because it is not chemically changing 

• Reversible processes 

Sublimation

Solid → gas 

• dry ice (solid CO2) sublimes from a solid directly to a gas 

Absorbed released

Deposition 

Gas→ solid 

• formation of frost in a freezer as the moisture in the air forms solid ice 

Released heat

Evaporation/ vaporization

Liquid→ gas 

• takes place only at the surface of the liquid 

• Can occur at temperatures below boiling point of the liquid 

Absorbed heat 

Melting 

Solid→ liquid

Absorbed heat

Freezing 

Liquid→ solid 

Released heat

Condensation 

Gas→ liquid 

Released heat 

Celsius scale

Based on the freezing point of water (0degrees Celsius)and the boiling point of water (100degrees Celsius)

Kelvin scale

An absolute temperature scale where the lowest possible value is 0K, known as absolute zero 

• at absolute zero, particles have zero kinetic energy

• The temperature in kelvin is directly  proportional to the average kinetic energy of the particles in the substance 

• does not have negative temperatures

Converting between K and degrees C

Add or subtract 273 depending on which scale you are converting from 

• C to K- add 273

• K to X- subtract 273

Heating and cooling curves

Changes of state graphed by producing a heating (or cooling) curve

Shows how the state of matter changes as heat is added; cooling curve would be the opposite, starting at a gas and ending with a solid with the temperature decreasing

• there are some points where the temperature remains constant because all the added heat is being used to overcome the intermolecular forces that act between the particles

Nucleons 

Located in the nucleus

Protons 

• Relative mass of 1 amu 

• Charge of +1

Neutrons 

• Relative mass of 1 amu

• No charge 

Electrons

Located in the electron cloud outside nucleus 

Relative mass of 1/2000

Negative charge 

Atomic number 

The number of protons in the nucleus of an atom 

Mass number/ nucleon number 

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

Nuclear symbol notation

Used to represent an element and can determine number of protons, neutrons, an electrons in an atom or ion

• Top number is mass number 

• Element symbol- center large 

• Bottom number is atomic number 

Ions 

A charged particle 

Has a charge as the number of protons do not equal the number of electrons 

Positive ions 

Formed when atoms lose electrons 

Has fewer electrons than protons 

Negative ions 

Formed when atoms gain electrons 

Has more electrons than protons 

Isotopes 

Atoms of the same elements that have different number of neutrons

Relative atomic mass 

The weighted average mass of an atom compared to 1/12 of the mass of the carbon-12 atom

• Calculated from the percent abundance and the masses of the isotopes of an atom:

  • mass # of each isotope is multiplied by its percent abundance, those values are added together and then divided by 100

Spectroscopy 

The study of interaction between matter and light 

emission spectra 

The range of frequencies or wavelengths of electromagnetic radiation emitted during an electron transition from a higher to a lower energy level 

Electron transitions

The movement of an electron between the energy levels in an atom, accompanied by the absorption or emission of energy

The electromagnetic (EM) spectrum 

Divided into 7 regions arranged in order of frequency, wavelength or energy

• Frequency and wavelength are inversely proportional

• The energy and frequency are directly proportional

• Radio waves have the lowest energy, lowest frequency, and longest wavelength

• Gamma rays have the highest energy, highest frequency and shortest wavelength

Wavelength λ

The distance between two crest in an an oscillating wave 

Units of distance (m)

Frequency( f)

The number of waves that pass a point in one second 

Units: hertz (Hz) or s^-1

Continuous spectrum

A spectrum that contains all the frequencies (or wavelengths) across a range of electromagnetic radiation

• our eyes see the continuous spectrum as white light

Emission line spectrum 

The range of frequencies or wavelengths of electromagnetic radiation emitted during an electron transition from a higher to a lower energy level 

lines get closer together (converge) at high energy, which corresponds to high frequency and short wavelength → the distance between the blue and violet lines on the hydrogen emission spectrum is smaller than the distance between the red line and light blue line

Continuous spectrum vs emission line spectrum 

• a continuous spectrum shows all the wavelengths or frequencies of visible light from red to violet 

• An emission line spectrum only shows specific wavelengths or frequencies of light. These are shown as colored lines on a black background 

Bohr model of the atom 

An atomic model that shows energy levels at fixed distances from the nucleus

Principal quantum number 

The main energy level occupied by electrons, assigned the letter n 

• n=1 is closest to the nucleus (known as ground state) 

• As the value of n increases, the distance from the nucleus and its energy increases

• N=1 has the lowest energy n= has the highest energy 

• Main energy levels converge at high energy 

Photons 

An elementary particle of discrete amounts of electromagnetic radiation 

Electrons transitioning between the energy levels 

 By either absorbing or emitting energy 

• the energy absorbed or emitted is in the form of photons (small packets of energy) 

• If an electron absorbs a discrete or an exact amount of energy, it will transition from a lower energy level to a higher energy level, for example from n = 2 to n = 3.

• The electron is now said to be in an excited state after absorbing (The excited state is unstable relative to the ground state.)

• The unstable electron emits the same amount of energy that it absorbed, and it transitions back down to n = 2.

• The amount of energy emitted by the electron in the transition from n = 3 to n = 2 corresponds to the wavelengths of visible light

• A line will be observed on the emission line spectrum

  ** the amount of energy emitted depends on the size of the transition (more energy is emitted for farther distance)

Sublevels

The smaller division of main energy levels, assigned to letters s,p,d,

• based on the shape of atomic orbitals

Recap: main energy levels are divided into sublevels which are made up of atomic orbitals

Atomic orbitals

A region of space where there is a high probability of finding an electron

• a single atomic orbital can hold a maximum of 2 electrons with specific orientations

S sublevel

• spherical

• Only consists of a single s atomic orbital so it I can hold maximum of 2 electrons

P sublevel

• composed of three p atomic orbitals

• Can hold a maximum of 6 electrons

D sublevel 

complex shape

Contains 5 d atomic orbitals 

F sublevel 

• Complex shape 

• contains 7 atomic orbitals 

Maximum number of electrons in an energy level

2n² where n is the principal energy level number

Isoelectronic 

Different ions of different elements with the same electron configuration

• ex. Mg²⁺ and Na⁺

Electron configuration

Shows the arrangement of electrons in their different levels around the nucleus of an atom

• For ions, take away or add the number of electrons at the end

Aufbau principle 

States that when adding electrons to an atom, the lower energy orbitals must be filled first 

Pauli exclusion principle 

States that an atomic orbital can only hold two electrons and they must have opposite spins 

Hund’s rule

States that when we have degenerate orbitals (orbitals of the same energy) then each orbital is filled with single electron before being double occupied

exceptions to orbital diagrams and electron configurations

Cr and Cu

• orbitals or sublevels want to be completely full OR half-full

• in Cr: one electron in 4s goes to 3d so BOTH are HALF-FULL→ more stable

Assumptions made in the ideal gas model

• the particles in a gas are in constant, random, straight-line motion 

• collisions between the particles are elastic(no energy lost)

• intermolecular forces between the particles are insignificant/negligible

• The average distance between gas molecules is much LARGER than the size of the molecules→ gas particles have negligible volume

• the average kinetic energy of the particles in a gas is directly proportional to the absolute temperature (in kelvin)

2 conditions under which real gases deviate the most from ideal gas behavior

• A high pressure

  • gas particles are closer together and are influenced by the forces of attraction

• At low temperatures

  • particles move less rapidly and therefore have less kinetic energy, so there is a greater opportunity for intermolecular forces between the particles to have an effect

ideal gas behavior

• low pressure

  • gas particles are far apart

  • forces of attraction are negligible

• high temperature

  • less opportunity for intermolecular forces between particles to have effect 

  • particles move more rapidly (more kinetic energy)

preparing a standard solution procedure steps

1. measure mass of solute using a mass balance 

2. Dissolve solute in a small volume of distilled water in a beaker and mixing with a stirring rod 

3. Transfer to volumetric flask via funnel

4. Rinse beaker, stirring rod and funnel with distilled water

5. Add distilled water to the volumetric flask until the bottom of the meniscus is at the 500 cm³ mark 

6. Add stopper. and invert to mix thoroughly

Molar mass of ideal gas formula

M=mRT/PV

SI units for ideal gas equations

Pressure- Pa

Temperature- K

Volume- m³

mole

the number of particles present in exactly 12 g of the carbon-12 isotope. this is equal to the Avagadro constant: 6.02 × 10²³ particles 

formula unit

the empirical formula for an ionic compound that represents the simplest ratio of ions making up the compound

SI

the International System of Units

relative formula mass

the mass of a compound relative to 1/12 the mass of an atom of carbon-12

• does not have any units

relative molecular mass

the mass of one molecule of a molecular compound, measured by comparison to 1/12 the mass of an atom of carbon-12 

• does not have units 

molar mass

the mass of one mole of a substance,

• units: g mol^-1

mole equation

n=m/M

• n-number of moles

• m-mass

• M- molar mass

Law of Definite Proportions

For any given compound, the ratio of constituent elements is fixed

• also known as law of definite composition or Proust’s law 

percent composition

the relative amount of each element in a compound

• will only give the simplest ratio of atoms in a compound (empirical formula)

empirical formula

the simplest ratio of atoms in a compound

molecular formula

the actual number of atoms in the compound

concentration

the number of particles or moles in a given volume, mol dm^-3

• concentrated and dilute are qualitative terms to describe concentration

concentration: mass per volume 

C=m/v 

• C- concentration: g dm^-3

• m-mass of solute:g

• v-total volume of solution dm³

concentration: in terms of moles per unit volume

C=n/v 

• C- concentration: mol dm^-3

• n-moles of solute: n

• v-total volume of solution: dm³

Factors that affect the number of particles of gas in a sealed container

• the volume of the container

  • larger the container, the more particles can fit in

• the temperature of the gas

  • greater temperatures mean more kinetic energy and therefore greater speeds the particles will travel

  • at low temperatures, the particles will move slowly, so more particles fit inside the container

  • at high temperatures, the particles need more space to move, therefore fewer particles can be accommodated

• the pressures of the gas 

  • at low pressures, there are fewer collisions with the sides of the container because there are fewer particles present

  • when more particles are present, the pressure will be higher

Volume, temperature, and gas determine the number of gaseous elementary particles that can fit in a container