Study of matter

I Structure  of matter

Pure substances and mixtures

Elements, allotropes, compounds, mixtures, separation of mixture, purification

Elements are different types of atoms, with unique sets of physical and chemical properties

• Identified by its atomic number Z = number of protons in its atoms

• Chemical symbols are unique one- or two-letter abbreviations

Isotope = atoms of the same element but with different amounts of neutrons, resulting in differing masses

• Similar / identical chemical behaviour, but may have different physical properties
  ( density, atom stability etc )

A way of distinguishing isotopes from their original elements is using an isotope notation

Allotrope = different forms of the same element, where the atoms are arranged in distinct ways.

• Differing physical and chemical properties (ie conductivity)

• Carbon : diamond, graphite, graphene, fullerenes

• Oxygen : diatomic oxygen, ozone

Molecule = 2 or more atoms held together by (usually) covalent bonds

• Smallest unit of a chemical compound, while still having the chemical properties

Compound = substance composed of two or more different elements that are chemically bonded together

• Ionic, covalent or metallic bonds

• Fixed ratios ( for each oxygen, there are two hydrogen in water )

• Need chemical methods to break their bonds ( ie. Electrolysis for water )

 All compounds are molecules, but not all molecules are compounds

Mixture = combination of two or more substances where each substance retains its own chemical properties

• Unlike chemical compounds, the compounds aren't chemically bonded together, which means that they can be separated by physical means

Homogeneous mixtures are uniform in composition and appearance throughout

• Solution = solute is dissolved in a solvent

• Alloy = a mixture of metals

• Gaseous mixture = mixtures of gas

• Colloid = particles of another substance are dispersed throughout another substance but not dissolved

Heterogeneous mixtures are not uniform and as such, the individual components can be seen and distinguished from one another

• Suspension = solid particles are dispersed throughout a liquid or gas, undissolved

• Colloid = can be considered heterogeneous as well if they consist of two distinct phases ( fog, whipped cream )

• Emulsion = may appear homogeneous, but unstable and might separate over time ( can be seen at microscopic levels that they are separated )

• Mechanical mixture = components physically distinct and can be easily separated ( cereal and milk )

Separation of mixture:

Homogeneous mixtures:

• Distillation = utilising differences in the boiling points of the components

• Evaporation = difference in boiling points or the volatility ( vaporisation ability )

• Fractional distillation = mixtures with close but not the same boiling points, may include multiple steps

• Chromatography = separates components based on their different ways of moving from stationary to mobile phases ( move at different rates through a tube )

• Ultrafiltration = semi-permeable membrane is used to separate particles based on size ( ie. water purification )

• Centrifugation = centrifugal force to separate components based on density
  ( spinning at high speeds to separate )

Heterogeneous mixtures:

• Filtration, sieving = differences in particle size

• Settling and decantation = settling of particles over time based on density

• Centrifugation = effective with heterogeneous mixtures as well, high speed spinning
  ( cream from milk )

• Magnetic separation = when one component is magnetic

• Chromatography (mentioned previously) different speeds of travel

Purification = separating substance from impurities and contaminants to obtain a pure form of a substance

States of Matter

Three states of matter (gas, liquid, and solid), changes of state

The main three different states of matter are determined by the arrangement and movement of its molecules or atoms

Solid = definite shape and volume

• Shape won't change unless physically altered

• Particles are closely packed in a fixed, orderly arrangement, vibrating but not moving freely

• Strong intermolecular forces hold the particles together, which gives solids their rigid structure

• Can be crystalline = regular repeating structure ( diamond, NaCl ),
  or amorphous = no organised pattern ( glass )

Liquid = definite volume, but takes the shape of its container

• Particles are close together but not in a fixed arrangement, they are able to slide past one another

• Particles move more freely, but still interact with one another

• Intermediate intermolecular forces hold particles together, allowing them to move around

• Surface tension due to the attraction between molecules at the surface

Gas = no definite volume or shape

• Expand to fill the shape and volume of their container ( easily compressible )

• Particles are far apart and not in any fixed arrangement

• Particles move rapidly and randomly, collisions with one another and the walls of the container and surroundings

• Weak intermolecular forces, allowing them to fill any available space

Changes of state occur due to changes in temperature and pressure, causing the general movement of particles to change

Energy absorbed:

Melting (solid to liquid), melting point

Vaporization (liquid to gas), reaching boiling point or evaporation

Sublimation (solid to gas),

Energy released:

Condensation (gas to liquid), condensation point

Freezing (liquid to solid), freezing point

Deposition (gas to solid)

Structure of the atom

Electron, proton, neutron, mass number, isotope

Atoms are basic units of matter, which are made up of subatomic particles: protons, neutrons and electrons

Electrons

• Found in the electron cloud located around the nucleus

• Electron cloud makes up most of the atoms volume

• Negative charge of 1- (or -1)

• Mass is about 0.0005u

• Determines the atom's chemical properties

Protons

• Found in the nucleus ( nucleon )

• Positive charge of 1+ (or +1)

• Mass is 1u

• The number of protons determines the element's identity and represented by its atomic number Z

Neutrons

• Found in the nucleus ( nucleon )

• No charge (0)

• Similar mass to protons, 1u

• Contribute to the mass of the atom and help stabilise the nucleus by offsetting repulsive forces between protons

 u = unified atomic mass unit (amu)

• Approx 1.665 540 * 10 ^(-27)kg

Atomic number Z = amount of protons in the nucleus of an atom, unique to each element

Mass number A = # of protons + # of neutrons

Average atomic mass ( atomic weight ) = elemet's weighted average mass of the atoms in a naturally occurring sample of the element

• Found below the chemical symbol on the periodic table

Isotope = atoms of the same element, but they have different amounts of neutrons, resulting in differing mass numbers

• The chemical properties are nearly identical since they still have the same amount of protons and neutrons

Electron configuration

Electron shell, properties of atoms, the periodic law, periodic table, valence electrons

Electron configuration = the arrangement of electrons in the electron cloud

• The configuration determines an element's reactivity, ionisation energy and placement in the periodic table

Electron shells = electrons are organised into different energy levels (= different shells)

Energy levels of electron shells increase with the distance from the nucleus

The electron shells are numbered from one, the closest being n=1, n=2, n=3 and so on

•  n = principle quantum numbers

• Letters can also be used:

  • K ( n=1)

  • L  ( n=2)

  • M (n=3)

  • N (n=4)

The maximum amount of electrons in an electron shell is calculated : 2n2 

The bohr model

• Nucleus is surrounded by rings, different energy levels or shells, which can hold different amounts of electrons

• Electrons cannot exist at energies between these levels

Lewis diagram = model where the nucleus is surrounded by just the valence shell (no ring)

Electron subshells =

• Four types: s,p,d,f

Orbitals

• Within each subshell, there are orbitals, where you'd most likely find the electrons

• Each orbital can hold up to 2 electrons

So with this information, lets use the M shell as an example

• It has 3 subshells (s, p, d)

• And in those subshells it has ( 1*2+3*2+5*2 ) 18 electrons

• Using the formula 2n2 (2*3^2 = 18 ) we can see that that amounts to the same amount of electrons found in the subshells.

Valence shell = the outermost shell of an atom

• Its electrons are valence electrons

• Most easily transferred or shared

Valence electrons determine most of an atom's chemical behaviours

• Elements in the same group have the same amount of valence electrons

  • Only for the main group elements (1A-2A, 3A-8A (1-2, 13-18) )

  • Except for helium, which has 2

= elements in the same group have the same amount of valence electrons, which means they have similar chemical properties and thus similar ions = similar roles in ionic compounds

Core electrons = non-valence electrons

Periodic law = when elements are arranged in order of increasing atomic number, their physical and chemical properties exhibit a periodic pattern

• Atomic radius decreases across a period (left to right) and increases down a group

• Ionisation energy increases across a period and decreases down a group

  • = energy required to remove an electron from an atom

• Electronegativity increases across a period and decreases down a group

  • = measure of an atom's ability to attract electrons in a chemical bond

The periodic table:

• Organised by increasing atomic number

• Elements are grouped by similar chemical properties

• Periods = horizontal rows, represent elements with the same number of electron shells

• Groups = vertical columns, contain elements with the same number of valence electrons and similar chemical properties

Ionic bonds

Ionic bond, ionic crystal, ionisation energy, electron affinity

Ionic bonds = based on electron transfer that happens between atoms, creating charges, which results in an electrostatic attraction that holds the ions together and forms the ionic bond

Elements in the same groups (same amount of valence electrons) have similar ions and roles in ionic compounds

Cation = lose a few electrons = positive ion

Anion = gain a few electrons = negative ion

The electronegativity difference between elements in an ionic bond has to be over 1.7

Ionic crystals are solid structures formed by a regular arrangement of ions held together by ionic bonds

• The repeating pattern forms a crystal lattice

• Unlike covalent bonds that form molecules and from there substances

Lattice energy = measure of the strength of the forces holding the ions together in the crystal lattice

• The higher, the more stable the ionic bond is

• Increases when the charges on the ions are greater or the size of the ions is smaller (can come closer together and experience stronger attraction

Ionic bonds can be broken by being dissolved in water or melting the substance, which provides enough energy to overcome the ionic attractions.

Typical properties of ionic bonds:

• High melting and boiling points

  • Strong ionic bonds require a significant amount of energy to break

• Hardness and brittleness

  • When force is applied and a layers of ions shift, ions with the same charge align and repel each other, causing the crystal to fracture

• Electrical conductivity 0 when solid

  • When melted or soluted, the ions are free to move, allowing the substance to conduct electricity (electrolytes)

• Soluble in water

  • Due to polarity in water; the positive and negative ends attract the oppositely charged ions, pulling them out of the lattice and into solution

• Crystal shape

Ionisation energy (ionisation potential, IE)

= the amount of energy required to remove an electron from an isolated atom or molecule

• Factors that affect the amount of energy required:

  • Atomic radius : the smaller the atom, the closer the electrons are to the nucleus and the higher the energy

  • Nuclear charge : attraction between nucleus and electrons is higher, which increases the needed energy

  • Electron shielding

Naming ionic bonds:

Formula names:

• Cation before anion "giver is respected"

• Net charge of 0!

• Lowest possible values

Naming cations:

Generally "element name + 'ion' "

• Hydrogen = > H+ , hydrogen ion

• Charge isn't mentioned, as it can be easily predicted

• ( some transition metals ie iron can be found with 2+ and 3+ charges, making it polyvalent ("many valued"), this is indicated by adding Roman numerals to the name, ie " iron(II)" and "iron(III)" )

Naming anions:

Suffix "ide"

• Hydrogen = > H- , hydride

• Chlorine = > Cl- , chloride

• Sulfur = > S^2- , sulfide

• Charge not mentioned when written

Electron affinity EA

= measure of the change of energy when an electron is added to a neutral atom

• Usually released energy

First electron affinity : energy change when one electron is added to a neutral atom

• Usually release energy ( exothermic process ) = negative electron affinity

Second electron affinity: energy change when a second electron is added to a negatively charged ion

• Requires energy, since adding an electron to an already negatively charged ion requires energy..

• Usually require energy ( endothermic process ) = positive electron affinity

So EA positivity ( + or - ) might seem opposite, but it describes the affected energy of the "system" of the atom or molecule, where losing an electron means lost energy (-), and gaining an electron means gained energy (+)

Applications and Importance

1. Chemical Reactivity

   • Elements with high (more negative) electron affinities tend to be more reactive nonmetals because they readily gain electrons.

2. Predicting Ionic Bonds

   • Electron affinity helps predict the likelihood of an atom forming an anion. Atoms with high electron affinities are more likely to gain electrons and form ionic bonds with metals.

3. Electron Affinity in Periodic Table

   • It aids in understanding and explaining the trends and reactivities of elements in the periodic table.

Notable exceptions

• Noble Gases: Generally have positive electron affinities because adding an electron would require placing it in a higher energy level, resulting in an endothermic process.

• Group 2 Elements (Alkaline Earth Metals): Typically have lower (less negative) electron affinities because the added electron would enter a higher energy p-orbital.

• Group 15 Elements (Nitrogen Family): Often have less negative electron affinities due to the extra electron entering a half-filled p-orbital, which involves electron-electron repulsion.

Salt = substance composed of cations and anions ionically bonded to each other

• Structure is solid, because ie NaCl, the Na+ attracts all Cl- ions that surround it

Metallic bonds:

Metallic bond, free electron, metallic crystal, malleability

• Electrostatic attraction between metal cations and delocalized electrons from the metals

• Form when metal atoms release some of their electrons into the "sea" of free electrons

The presence of free, delocalized electrons from the metal atoms explain some of the main structural features of metallic structures:

• Conductivity: the free movement of electrons makes metals excellent conductors of electricity and heat

• Malleability: when force is applied to the structure, the sliding of the structure won't cause it to break since the electrons will still surround and prevent atoms with similar charges from repelling each other

A metallic crystal is a solid structure composed of a regular, repeating arrangement of metal cations surrounded by a sea of delocalized electrons.

• Structure: The metal cations form a lattice, a repeating 3D pattern that extends in all directions.

• “Delocalised electrons are almost like a fluid filling the spaces between fixed metal atoms”

The melting and boiling points of metals depend on the strength of the metallic bonds

• Generally high, since it requires a lot of energy to overcome the strong attraction between metal ions and delocalised electrons

Factors affecting metallic bond strength:

• Number of delocalised electrons; more electrons available to participate in the bonding increases bond strength

• Atomic size; smaller atoms result in stronger metallic bonds (nuclei are found closer to the sea of delocalised electrons)

• Charge of metal ions; the higher positive charges on the metal ions, the stronger the bond.

Alloy = mixture of two or more elements, where at least one is a metal

• Designed to improve the properties of pure metals ( strength, hardness, resistance to corrosion )

Naming metals

Pure metals are named after the element themselves.

Iron (Fe)

Copper (Cu)

Silver (Ag)

Alloys either have a special name or simply state the mixture.

Brass = copper and zinc

Bronze = copper and tin

Steel = iron with carbon

Stainless steel = steel with chromium and nickel

Aluminum-magnesium alloy

Copper-nickel alloy

Covalent bonds

Covalent bond, coordinate bond, crystal of covalent bond, molecular crystals, polar nature of bond, electronegativity

= chemical bond formed by the sharing of electron pairs between atoms

• Atoms want to attain the electron configuration of a noble gas, resulting in a stable molecule

• The sharing can be equal or unequal, depending on the electronegativity of the atoms involved

Molecule = group of two or more atoms covalently bonded together

Diatomic elements = elements found as covalently-bonded pairs in nature

• Considered molecules, not compounds

Covalent network solid = some covalent substances are bonded together in a network instead of just one molecule

• ie. silica

Single covalent bond = 1 shared pair of electrons, 2 electrons total

Double covalent bond = 2 shared pairs of electrons, 4 electrons total 

Coordinate bond (dative covalent bond) = covalent bond where both electrons in the shared bond come from the same atom

Crystal of covalent bond = solid, where the atoms are connected by covalent bonds

• Diamond is a crystal of covalent bonds, carbon atoms in a tetrahedral structure

Electronegativity = measure of an atom's ability to attract and hold onto electrons in a covalent bond

• Increases across a period from left to right, decreases down a group

• ="how much they want to "steal" electrons

Electronegativity is measured with the Pauling scale

• Ranges typically from 0.7 to 4.0

• A relative scale

Polarity of covalent bonds: polar = "unequal sharing" of the electrons,

• Nonpolar electronegativity difference is between 0 - 0.4

• Polar electronegativity difference between 0.5 - 1.7

• Ionic >1.7

Properties of covalent bonds:

• Polarity

• Low melting and boiling points, due to the intermolecular forces being generally weaker than the bonds themselves

• Polar covalent compounds are soluble in polar solvents (water)

• Nonpolar covalent compounds are soluble in nonpolar solvents (oil)

• Not conductive

• Molecular structures; each molecule is their own "unit" instead of a continuous lattice seen in ionic or metallic compounds.

Naming covalent bonds

First element name is usually just the element

Second element name takes the root of the element’s name and adds the suffix -ide.

Prefixes indicate the number of atoms of element present in the compound

Prefixes:

Mono (usually omitted from the first element)

Di

Tri

Tetra

Penta

Hexa

Naming elements and filling in reactions practise

Intermolecular forces

van der Waals force, hydrogen bond

Intramolecular forces = forces that hold together atoms within a molecule

• Very strong compared to intermolecular forces

• Responsible for the chemical properties of a substance

• Covalent, ionic and metallic bonds

Intermolecular forces = forces of attraction or repulsion between molecules

• Determines the physical properties of a substance (melting and boiling point, vapor pressure, solubility)

• Van der waals forces, hydrogen bonds

van der Waals force 

= weak intermolecular forces that occur between molecules or atoms, not involving chemical bonds, but instead arising from temporary attractions due to electron movement.

• There are three types on Van der Waals forces:

1. London dispersion forces (dispersion forces)

   • Weak attractions due to temporary fluctuations in the electron distribution within atoms or molecules (regardless of polarity)

   • Increases with the size and shape of the molecule, because there's a larger surface area where temporary dipoles can form

2. Dipole-dipole interactions

   • Interactions between polar molecules, due to the electrostatic attractions, between polar molecules

   • Generally stronger than LDFs but weaker than hydrogen bonds

3. Dipole-induced dipole interactions

   • A polar molecule induces a temporary dipole in a neighbouring nonpolar molecule

     • A permanent polar molecule can distort the electron cloud of a nearby nonpolar molecule, inducing a temporary dipole

   • Weaker than normal dipole-dipole, generally stronger than LDFs

Hydrogen bond

= strong type of dipole-dipole interaction

• A hydrogen atom is covalently bonded to a highly electronegative atom (most commonly nitrogen (N), oxygen (O) or fluorine (F)) which then interacts with a lone pair of electrons on another electronegative atom in a nearby molecule

  • Lone pair of electrons = pairs of valence electrons not shared with another atom and are not thus involved in bonding

    • Provide a region of negative charge

Molecular polarity

1. Electronegativity (> 0.5 - 1.7 difference generally)

2. Molecular geometry (asymmetry needed (not linear for example))

3. Lone pairs (negative zones)

Chemical bonds and properties of substances summary:

Amount of substance

Atomic weight, molecular weight, formula weight, amount of substance, molar concentration, mass percent concentration, molarity 

Atomic weight  (average atomic mass)

= weighted average of different isotopes of an element

Molecular weight (molar mass)

= element's atomic weight expressed in g/mol

• Covalent compounds

Formula weight (can be called molar mass as well)

= essentially calculated the same as molar mass, but used for ionic compounds that don't form molecules

Amount of substance

• Measured in moles, one mole is defined as particles (Avogadro's number)

Molar concentration (molarity)

= number of moles of solute per liter of solution (mol/L)

Mass percent concentration (weight percent)

= ratio of the mass of the solute to the total mass of the solution in %

Chemical formulas

Molecular formula, ion formula, electron formula (Lewis structures), structural formula, compositional formula (empirical formula)

Molecular formula = represents the exact number of atoms of each element in a molecule

• Shows total composition but not the structure of the molecule

Ion formula = shows the charges of individual ions

Electron formula (Lewis structures) = depict the arrangement of valence electrons around atoms

Structural formula = illustrates how atoms are arranged and bonded within the molecule

Compositional formula (Empirical formula) = provides the simplest whole-number ratio of atoms in a compound instead of the exact number of atoms. Useful when analyzing chemical compositions and reactions where only the relative proportions are needed

Recap:

• Molecular Formula gives the actual number of atoms, while Empirical Formula simplifies the ratio.

• Lewis Structures provide insights into electron arrangements and bonding.

• Structural Formula focuses on the physical arrangement of bonds.