Preliminary Chemistry

Mixtures

A mixture is a combination of two or more pure substances (containing one type of molecule) - the mixture itself is impure.

Heterogeneous Mixtures

Two or more substances intermingle but ​remain physically separate​.

E.g dirt and sand, oil and water, salt and baking soda

A suspension is a specific type of heterogeneous mixture where particles settle at the bottom.

Homogeneous mixture

Two or more substances have merged into a ​uniform phase​. There are no borders between the substances, but they are not chemically bonded. The ​physical properties​ of each ingredient can be exploited to separate them.

E.g. Saltwater, Copper Sulfate solution, aqueous solutions

Elements vs Compounds

Elements are pure substances that cannot be chemically or physically decomposed.

Compounds are pure substances that are chemical combinations of two or more different elements - they can be decomposed.

Periods

The rows of the periodic table. Each period corresponds to the number of electron shells​ of that elements.

Groups

The columns of the periodic table. Elements in the ​same group​ share similar chemical properties, as they have the ​same number of valence electrons​.

Periodic properties

Atomic Radius, Ionisation Energy and Electronegativity

Atomic Radius

Half the distance between the centers of two atoms of an element that are touching

Going ​left → right​ across a period, atoms have more protons but the same amount of electron shells. Thus, Electrons are attracted to the nucleus more strongly, and the atomic radius decreases

Going ​up → down​ the group, atoms have more electron shells, which not only put the valence electrons further away, but the inner electrons also repel (or shield) the valence electrons from the nucleus's attraction, so the ​atomic radius increases

Atomic radius affects all the other properties​ - i.e. it's easier for an atom with a greater atomic radius to let go of an electron, as greater ionisation energy

Ionisation Energy

​The energy required to remove one valence electron from a gaseous atom.

The more strongly bound to the nucleus electrons are, the more ionisation energy is required to remove them

Smaller atomic radii mean stronger bound electrons, so​ ionisation energy increases as atomic radius decreases

A low first ionisation energy indicates that an element is a metal, while a high first ionisation energy indicates that it is a nonmetal

Electronegativity

The measure of the ability of an atom to attract electrons for chemical bonding (measured in ​Pauling units​).

When an atom has a smaller atomic radius and the atom can easily pull external electrons into it. Thus, ​as atomic radius decreases, electronegativity increases

FON = most electronegative

Periodic Trends

Moving Left → Right (Periods):

Ionisation Energy Increases

Electronegativity Increases

Atomic Radius Decreases

Moving Up → Down (Groups):

Ionisation Energy Decreases

Electronegativity Decreases

Atomic Radius Increases

Isotopes

Elements with the same number of protons and electrons but a different number of neutrons.

e.g. a Hydrogen atom can have 0, 1 or 2 neutrons, but it is still hydrogen.

Isotope Trends

Isotopes with atomic number > 82 are all unstable.

Isotopes with 1:1 proton-neutron ratio are much more likely to be stable.

Atoms with an even number of protons and neutrons are more likely to be stable.

Mass spectrometer

Device that uses electromagnetic fields to sort the isotopes present in a substance by atomic mass, which then allows us to see how abundant each isotope is.

Half life

Measure of the time it takes for half the atoms in that substance to decay. Half-lives can range from seconds to billions of years, and can be represented as a logarithmic graph.

Types of radiation

Alpha, Beta, Gamma

Alpha Decay

Occurs when an atoms mass is too much therefore emits an alpha particle a Helium particle (2 protons, 2 neutrons)

Atomic number decreased by 2, mass decreased by 4.

For example, uranium-238 transforms into thorium-234.

Beta Decay

Occurs when an atom has too many neutrons therefore emits a beta particle - proton changes into a neutron and a positron; the positron is emitted.

Atomic number decreases by 1, mass is conserved.

For example, thorium-234 (90) becomes protactinium-324 (91)

Gamma Radiation

Occurs when an atom emits gamma rays. It usually occurs after alpha or beta decay, where the nucleus is still excited after decaying. The excited nucleus then releases gamma ray photons to become more stable. Gamma Radiation technically isn't a type of decay because only energy is released.

Penetrative power of each radiation

Alpha: Few centimetres in air

Beta: Few millimetres of aluminium

Gamma: Many centimetres of lead

Bohr Model

1. Electrons are particles that occupy ​fixed orbits​ around the nucleus - called stationary orbits

2. Each orbit has an ​energy level​ associated with it

3. Energy is absorbed when an electron jumps from a lower orbit to a higher one and energy is emitted when an electron falls from a higher to a lower orbit

4. Electrons cannot exist between the energy levels/orbits

Energy Levels and Electron Configuration

Electrons do not orbit the nucleus in fixed orbits, they move around in energy levels called

orbitals - which are regions around the atom where the electron is very likely to be found at any given time​.

Aufbau Principle

States that in order to configure electrons in an atom, electrons are added to the lowest energy level until it is filled. Then electrons are added to the next energy level until that is filled, etc.

spdf notation

The principal shells are named in increasing order of energy as 1,2,3, etc, while the subshells are named s, p, d, f. Each subshell can hold a maximum amount of electrons, depending on the orbitals that it contains.

For example titanium which has 22 electrons​

1st diagonal: 1s (add 2 electrons) → 20 electrons left

2nd diagonal: 2s (2 electrons) → 18 electrons left

3rd diagonal: 2p (6 electrons), 3s (2 electrons) → 10 electrons left

4th diagonal: 3p (6 electrons), 4s (2 electrons) → 2 electrons left

5th diagonal: 3d (can hold 10 electrons, but add 2) → 0 electrons left

1s2​ ​, 2s2​ ​, 2p6​ ​, 3s2​ ​, 3p6​ ​, 3d2​ ​, 4s2​ ​

Bonding

In order to gain a full shell or noble gas configuration, atoms will remove/gain electrons to make itself more stable. ​Electronegativity​ is the main mechanism behind bonding.

Ionic Bonds

Involve a Metal donating electrons to a Nonmetal atom.

The ​metal​ atom becomes ​cation) (+) after losing electrons while ​nonmetal ​atom becomes ​anion (-) after gaining electrons, creating ionic compound.

E.g. sodium chloride where sodium (cation) and chlorine (anion).

Covalent Bonds

Involve a nonmetal sharing electrons with another nonmetal atom. They bond as these electrons are shared to form a covalent or molecular compound.

E.g. Hydrogen bonds with Oxygen, two hydrogen atoms each share an electron with an Oxygen atom and form Water (H​2O​)

Polarity of Covalent Bonds

Electrons can be shared equally in covalent bonds, however not often. Occurs when one atom has ​higher electronegativity (wants the electrons more).

Often occur between atoms that are the same, such as two Cl atoms or H atoms.​

In ​Polar Covalent Bonds​, electrons​ NOT shared equally ​between the atoms, because the electronegativity of the atoms is different.

Atom with higher electronegativity = slightly negatively charged (δ−)

Atom with lower electronegativity = slightly positively charged ( δ+).​

Bond is partially ionic​.

The higher the electronegativity difference, the more _______ the bond is

ionic

The electronegativity of an element can generally be predicted by its periodicity - so two elements such as Hydrogen and Bromine which are located far from each other would make a very polar covalent bond.

Lewis Dot Diagram example

Carbon dioxide

Intermolecular forces ​

Are the forces that exist ​between m​olecules. These forces can also arise from electromagnetic attractions holding molecules together.

Intra vs Intermolecular forces

Intramolecular bonds are much stronger than intermolecular forces - breaking apart the atoms inside a molecule requires a lot more energy than breaking molecules away from each other.

The larger atomic radius, the ... its intermolecular forces

The more electrons an atom has (i.e larger atomic radius), the stronger it's intermolecular forces.

If it is a solid, it has stronger intermolecular forces than a substance that is a gas at room temp.

Dipole-Dipole Interactions​

When multiple molecules are in polar covalent or ionic bonds, one side of the molecule is slightly positive, and one side is slightly negative. These sides are known as dipoles. When the negative dipole of one molecule attracts to the positive dipole of another, they bond. This is the strongest intermolecular force.

Hydrogen Bonds

​a special kind of dipole-dipole interaction that occurs between hydrogen atoms bonded to either FON atoms. The positive dipole of Hydrogen is attracted to the negative dipole of fluorine/oxygen/nitrogen. Due to the large electronegativity difference, this is the strongest version of dipole-dipole interactions.

Dispersion Forces

The weakest intermolecular force. They intrinsically exist between all molecules (doesn't have to be ionic/polar) and occur due to the movement of electrons forming temporary dipoles. The more electrons a molecule has, the stronger the dispersion forces.

Physical Properties due to bonding and intermolecular forces

Polar covalent/Ionic compounds have strong dipole-dipole interactions between them, while Nonpolar compounds have only weak dispersion forces. The stronger a bond is, the more energy is required to break it, so ionic and polar compounds tend to have higher melting and boiling points than nonpolar compounds.

E.g. Water, a polar covalent compound, boils at 100°C while sodium chloride, an ionic compound, has a much higher boiling point at 1413°C.

Chemical structures

Ionic network, Covalent network, Covalent Molecular and Metallic.

Ionic Network

Ionic compounds​ form lattice structures, where each cation is surrounded by 6 anions, and each anion is surrounded by 6 cations - forming something like a cube like NaCl does.

Ionic solids have poor electrical conductivity - they have no free electrons, because they are locked into the oppositely charged atoms, and the ions are stuck in a lattice structure.

Covalent molecular

Discrete molecules​, with ​weak intermolecular forces between each molecule​.

Covalent Molecular substances do not conduct electricity because there are no ions or free floating electrons.

Covalent Network

Can be 2D and 3D - most prominent example is Graphite (2D Network) vs Diamond (3D Network). They are both different structures of the same element and have very different properties because of their different structures.

Metallic

A lattice of positive metal ions are held together by delocalised electrons (electron cloud/ electron sea) in metallic lattices.

The delocalised electrons carry electricity very well, making metals great conductors of electricity.

They also carry heat energy in the form of kinetic energy by moving around and hitting the colder parts of the lattice, making metals great conductors of heat

Delocalised electrons are also responsible for the malleability/ductility of metals

Valence shell electron pair repulsion theory - CO2 AND H2O

The shape of a molecule depends on the electron configurations of its atoms - theory.

E.g. CO​2​ has its nonbonding electrons distributed equally on opposite sides (it is nonpolar), so has a linear structure, since the repulsion is equal on either side.

H​2O​ has it's nonbonding electrons on top of oxygen, so they repel the hydrogen atoms downwards. As the hydrogen atoms come closer together, they also repel each other, creating a bent molecular shape

Stoichiometry

Since the Law of Conservation of Mass states that matter cannot be created or destroyed, the atoms before and after the reaction must be the same.

Thus, Chemical Equations can be balanced by adding stoichiometric coefficients​ in front of the molecules.

E.g.​2Na(s)+Cl​2​ (g)→2NaCl(s)

This means that for every two separate Sodium atoms (total mass: 45.98 amu) and diatomic Chlorine molecule (70.90 amu) that react, therefore ratio is 2:1.

Molar Mass Formula

n = m/MM

Number of Moles = Mass(g) ÷ Molar Mass(g/mol)

Concentration Formula

n = CV

(Number of Moles = Concentration of Solution (moles/litre) * Volume (litres))

Limiting reagents

Substance that is totally consumed when the chemical reaction is complete. The amount of product formed is limited by this reagent, since the reaction cannot continue without it, and will leave an excess of another substance.

Concentration

C​1V1​ = C​2V​2

*​Convert concentrations to mol/L if required in calculations

Avogadro's Gas Law

When measured at the same temperature and pressure, equal volumes of different gases contain the same number of molecules.

V​1/​ n​1​ = V​2/​ n​2

Charles' Law

The volume of a gas is directly proportional to its temperature when measured at a constant pressure. Convert celsius to kelvin.

V​1/​ T​1​ = V​2/​ T​2

Gay-Lussac's Pressure and Temperature Law

The ​pressure​ exerted by a gas is​ directly proportional​ ​to​ its ​temperature ​(in Kelvin), when

Volume is constant. - Kevlin instead of celsius

P​1/​ T​1​ = P​2/​ T​2

Boyle's Law

The ​pressure​ exerted by a gas is ​inversely proportional to​ its ​volume ​when measured at a constant temperature.

P​1V​1​ = P​2V​​2

Kelvin to celsius

celsius + 273.15 = kelvin

Combined Gas Law

P1V1/T1 = P2V2/T2

Temp always in Kelvin

STP = standard temp (273 K)

Standard Solutions

Known and fixed concentration.

Ideal Gas Law

Theoretical model of a gas that obeys the gas laws exactly. It follows certain assumptions:

- gas molecules move randomly in straight lines

- pressure is due to collisions between the walls of the container and the molecules

- no intermolecular forces between gas molecules

- gas molecules themselves take up no volume

NOT ACTUALLY TRUE FOR GASES

PV = nRT

R is the universal gas constant: 0.082 K⋅mol/L⋅atm:

Electrolysis

Application of an electric current to decompose a compound Consider: 2H​2O​ → 2H​2​ + O​2

Requires much more energy than boiling water, since it's breaking covalent bonds instead of only the intermolecular bonds. It is also a chemical reaction as bonds are being broken and reformed, while boiling water is a physical change.

Chemical Reactions

When chemical bonds are broken, rearranged and established to form new substances. They involve reactants turning into products.

Indicators of a chemical reaction

1. Bubbles

2. Colour change

3. Change in Energy (change in temperature)

4. Appearance of a solid (due to precipitation)

5. Disappearance of a solid (not due to dissolving)

6. Formation of new substance ← this is the only thing that guarantees a chemical reaction occurring,

the rest are ​indicators

Exothermic reaction

The ​reactants have more energy than the products

Energy is released​ from bonds being formed and goes into the surroundings, usually causing the ​temperature to rise.

E.g. Combustion

Endothermic reaction

Energy is taken from the surroundings and the ​reactants have less energy than the products. Because ​energy is being used ​to break bonds, endothermic reactions usually cause the ​temperature to drop.

E.g. Decomposition

Synthesis Reactions (A + B → AB)

Formation of compounds by combining simpler substances such as elements

Exothermic (energy is given off when bonds are broken)

E.g. 2Na + Cl​2​ → 2NaCl

Displacement Reactions - AB + CD → AD + CB

When salts dissolve in water, the bonds between the cation and anion separate inside the solution.

This swapping of ions when two salt solutions are mixed.

Reactants have to be soluble

E.g. Na​2S​ + 2HCl → 2NaCl + H​2S​

Precipitation Reactions - Soluble Salt (aq) + Soluble Salt (aq) → Precipitate (s) + Soluble Salt (aq)

New insoluble salt is formed and falls to the bottom if heavy enough, new solid particles formed are called the ​precipitate.

E.g. ​AgNO​3​ + NaCl → AgCl + NaNO​3

Decomposition Reactions (AB → A + B)

when atoms of a compound are separated to form two or more products.

E.g. NaCl→Na+Cl

Neutralisation reaction

Acid + Base → Salt + Water

E.g. HCl + NaOH → NaCl + H20

Acid and metal reaction

Acid + Metal → Salt + Hydrogen gas

E.g. Mg + 2HCl → MgCl​2​ + H​2

Metal Carbonate + Acid

Metal Carbonate + Acid → Salt + Carbon Dioxide + Water

E.g. 2HCl + Na2CO3 → CO2 + H2O + 2NaCl

Metal Oxide + Acid

Metal Oxide + Acid → Salt + Water

E.g. CuO+2HCl→CuCl2​ ​+ H2​ ​O

Combustion reaction

Hydrocarbon + Oxygen → Carbon Dioxide + Water

Incomplete forms → Carbon Monoxide + Carbon + Water

E.g. Complete Methane CH​4 + 2O2​ ​ → CO2​ ​+ 2H2O

E.g. Incomplete for Ethane C​2​H6 ​+ 2O2​ ​→CO + C + 3H2​O​

Aboriginal Detoxifying Poisonous Foods

Detoxifying cycad macrozamia was exploited as an important food source in spite of its being highly toxic and carcinogenic.

1. Aborigines cooked the seeds in a fire pit to denature the toxic compounds - denaturing refers to using heat to destroy its characteristic properties by disrupting its molecular structure.

High temperatures break the bonds between the molecules in the structures removing their ability to disrupt liver and neurological functions.

2. Slice the nuts finely and put them in a basket to soak in running water for a few days, dissolve any leftover amounts of toxic compounds remained. Cycasin especially can be hydrolyzed (it breaks down as it reacts with water).

3. Then cook it again before eating.

Reactivity Series of Metals

Polly Smith Cut My Ancient Zebra Into Large Chunks Making Super Glue.

K, Na, Ca, Mg, Al, Zn, Fe, Pb, Cu, Hg, Ag, Au

Metal Reactivity in terms of Periodic Trends

Since metals want to give away electrons, the more reactive metals give away electrons more easily.

Thus, metals with larger atomic radius (and therefore lower ionisation energies, and lower electronegativity) are more reactive.

Rate of Reaction

Measure of how quickly reactants are turned into products - it is represented by a change in the concentration of reacts/products with time.

Links to collision theory (must collide with correct orientation and energy)

Way to Increase RoR

Increasing Temperature​: gives the particles kinetic energy for collision

Increasing Concentration:​ more particles of reactant in the same space, so collisions are more likely

Increase the Surface Area​: more of one reactant is exposed to the other, so more collisions

Add a Catalyst​: a substance which speeds up the reaction by decreasing the activation energy required

Increasing Pressure​: which pushes particles closer together, as increasing pressure decreases volume

Activation Energy

Minimum amount of energy that is required for a set of reactants to begin reacting.

When met the existing bonds break and reactants collide at high enough speeds

Activation Energy Curve

The activation energy is the peak of the curve. A catalyst (such as an enzyme) decreases the activation energy required - blue

Energy Changes in Chemical Reactions

All chemical reactions have some energy changes occurring, since bonds are being broken and formed.

● The ​formation of bonds produces energy

● The​ breaking of bonds consumes energy

These energy changes follow the law of conservation of energy - the things storing Chemical Potential Energy are:

● Chemical Bonds

● Intermolecular Forces

● Repulsion between Electrons

● Repulsion between Nuclei

● Kinetic Energy of particles

Thus, the energy of the system isn't the same before and after the reaction - some reactions release energy, while some consume energy (i.e. exothermic and endothermic reactions). This is dependent on what and how many bonds are being broken/formed.

Enthalpy

Enthalpy is defined as the energy content of a system.

It is very difficult to know the exact enthalpy of a substance, so the enthalpy change is measured/calculated instead.

Measured as J/mol.

ΔH = -Q/n

+ve enthalpy = endothermic

-ve enthalpy = exothermic

Hess' Law

Hess's Law ​states that the total enthalpy change for a chemical reaction does not depend on

the pathway that it takes - it only depends on the initial and final states.

The total enthalpy change for a reaction carried out in multiple steps is also equal to the sum of the enthalpy changes of each individual step.

Energy Profile Diagram

Energy Profile Diagrams are representations of the energy (or enthalpy) level of the reaction as the reaction progresses.

For example, it might start at one point, then gain energy as bonds are broken, then release energy as new bonds are formed, until the final products have an net lower enthalpy than the initial reactants. This shows that the reaction has overall released some energy, making it exothermic

If the ​Energy of Products < Energy of Reactants​, it is ​exothermic​ (because that energy is released as heat), and ΔH < O.

If the ​Energy of Products > Energy of Reactants​, it is ​endothermic​ (because additional energy is consumed), and ΔH > O.

Also note that the Activation Energy in an exothermic reaction is generally lower than an endothermic reaction, because endothermic reactions are actually consuming energy.

Bond Energy

Bond Energy/Enthalpy​ refers to the amount of energy required to break the bonds of 1 mol

of a substance into its constituent atoms under STP.

E.g. Given that an N-N triple bond has a bond energy of 945 kj/mol, H-H bond has a bond energy of 436 kj/mol, and N-H bond has a bond energy of 391 kj/mol, calculate the enthalpy change of the following reaction:

N​2​ + 3H2​ ​ → 2NH​3

Energy consumed to break bonds = 1 x 945 + 3 x 436 = 2253

Energy released by new bonds forming = 6 x 391 = 2346

ΔH = ​[phone] = -93 kj/mol ​(exothermic)

Therefore,

ΔH = ∑ ΔH​(bonds broken)​ - ∑ ΔH​(bonds formed)

Standard Enthalpy of Formation

The standard enthalpy of formation of a compound is the ​change of enthalpy during the formation of 1 mole of the substance ​from its constituent elements, with all substances in their standard states.

Giventhestandardenthalpiesofformation(ΔH​f)​ of different substances, the change in enthalpy of a reaction can be calculated:

ΔH​(Reaction)​ = ∑ΔH​(Products)−​ ∑ΔH​(Reactants) Remember to multiply the

Stepped Reactions

A stepped reaction is one that occurs in multiple steps. When given information about the enthalpy changes for each step, the total enthalpy change of the reaction can be calculated.

E.g. Calculate ΔH for the following reaction: 2CO + O​2​ → 2CO​2 2C + O​2​ → 2CO, ΔH = -221.0 kJ/mol

C + O​2​ → CO​2,​ ΔH = -393.5 kJ/mol

1. Manipulate the given reactions to resemble the actual reaction

a. 2CO → 2C + O​2​ (equation flipped so flip ΔH, ΔH = 221.0 kJ/mol)

b. 2C + 2O​2​ → 2CO​2​ (equation x2 so ΔH x2 as well, ΔH = -787 kJ/mol)

2. Add the manipulated reactions

a. 2CO+2C+2O​2​ →2C+O​2​ +2CO​2

b. ΔH is added too, so 221 - 787 = -566 kJ/mol = ΔH of this reaction

3. Cancel out the same compounds on both sides, then check if the equation matches the original

a. Cancel out 2C and one O​2

b. 2CO + O​2​ → 2CO​2​ is left, this matches the initial reaction

c. So the answer is ΔH = -566 kJ/mol

Calorimetry

A calorimeter is a device used to measure the quantity of heat flow in a chemical reaction. The type shown in the diagram below is a "coffee cup calorimeter", used to measure neutralisation and similar reactions

● It consists of a covered container surrounded by insulation (such as styrofoam), to minimise heat getting in from the surroundings, because it only wants to measure the heat flow in the chemical reaction.

● A thermometer measures the temperature of the system as the reactants are stirred.

Bomb calorimeters

are a more effective alternative to accurately determine enthalpy changes in combustion reactions. This encompasses a copper 'bomb' containing measured amounts of reagents, which is submerged in a calorimeter of water. The chemical is ignited,

and the heat released warms the water, making use of the ​specific heat of water​ to determine change in energy within the system.

​specific heat formula​

can be used to calculate the heat/enthalpy change:

Q = mc∆T

Q = heat energy (Joules, J)

m = mass of a substance (g - this is can also be kg depending on the units used in specific heat capacity)

c = specific heat capacity (units J/g/K)

∆T = change in temperature (Kelvins, K, or Celsius, C - this would be the same)

The specific heat capacity of a substance is the amount of energy needed to increase 1 gram of that substance by 10​ ​C.

● Note in bomb calorimeters the mass and specific heat capacity of water should be used as m and c

Molar Heat of Combustion

refers to the energy released when 1 mole of a substance undergoes complete combustion under STP - it is measured in kJ/mol. This is always a positive​ value, as combustion is an ​exothermic r​ eaction, and energy is released.

● The ​Molar Enthalpy of Combustion​ is always a ​negative​ value, because it is the change in energy of the system per mole of substance that undergoes combustion.

One method used to determine enthalpy of combustion of​ liquid fuels​ (eg. ethanol, methanol) is to burn the fuel in a spirit burner, and using the heat from the combustion of the ethanol to heat a measured volume of water, held in a double-walled copper calorimeter.

● Copper ensures efficient heat transfer, as it is a conductor, but it is often double-walled​ to prevent the heat exiting the calorimeter

● Make sure the flame is touching the calorimeter

● Leave only a hole at the top for the thermometer

Calculating Molar Heat of Combustion

The total amount of heat released by a substance is the number of moles of that substance x the molar heat of combustion:

Q = mcΔT

Q = nH​C​ (when there is 100% heat transfer)

Enthalpy of Combustion:

use ∆H = -mc∆T/n NOTE: m = mass of ​water​ used, not of the fuel

Unfortunately, the results of this technique are not very accurate; a great deal of heat can be lost in the surroundings, and the combustion may be incomplete (indicated by the deposition of black soot on the bottom of the water container).

Exothermic energy profile diagram

Endothermic energy profile diagram

Validity

1. Controlled variables

2. Does it answer the aim?

3. Not repeated unreliable therefore invalid is it accurate

Covalent molecular

Covalently bonded molecules held together by weak intermolecular forces

Why does calorimeter experiment not produce results close to the published values?

As heat is lost to the surrounding air, to the container holding water. Sometimes combustion is incomplete.

Percentage Error formula

energy absorbed during the experiment / actual value (secondary source) x 100

Enthalpy Change

Nonmetals

Nonmetals​ are (usually) good insulators of heat and electricity, are brittle; usually dull many of the elemental nonmetals are gases at room temperature, while others are liquids and others are solids.

Metalloids

Metalloids​ have properties of both metals and nonmetals, and can be made to conduct electricity in some circumstances.