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Chemistry (Unit 1)
revision for mid year exam
| Term | Definition |
|---|---|
| Atom | > smallest unit of matter. |
| Proton | > positive charge in nucleus. - atomic number |
| Neutron | > neutral/no charge in nucleus. |
| Electron | > negatively charged, in clouds/shells. |
| Molecule | > when two or more of the SAME atoms bond together. eg. O₂ |
| Compound | > when two or more DIFFERENT atoms bond together eg. H₂O |
| Mass number | > number of protons and neutrons in the nucleus of an atom. - mass number = number of protons + number of neutrons |
| Isotopes | > atoms of the same element with different mass numbers. - different amount of neutrons = different mass numbers. - identical chemical properties, different physical properties. |
| Elements | > pure substance containing only a single type of atom. |
| Atomic radius | - decreases across a period (more protons as you go across cause valence electrons to be more attracted to nucleus). - increases down a group (new shell is filled). |
| Valence electrons | > electrons in the most outer shell/energy level. - same in each group (eg. Group 1 has 1 valence electron.) |
| Electronegativity | > the tendency of an element to attract electrons towards itself. - higher core charge = greater electronegativity > increases across a period, decreases down a group. |
| First ionisation energy | > the amount of energy than an atom must absorb in order to release/lose ONE electron (to become a cation). - higher core charge = greater first ionisation energy > increases across a period, decreases down a group |
| Core charge | > the attraction between the valence electrons and the nucleus. - Core charge = number of protons in the nucleus – number of total inner-shell electrons |
| Critical Elements | > elements that are vital for industry and technology play an important role in the develoapment of the human race. - demand is increasing at a rate that is not sustainable. |
| Electron configuration | > electrons fill the subshells from the lowest energy subshell to the highest. - (in order of increasing energy level) 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, and 7p. |
| Covalent bond | > involve 2 non-metal atoms SHARING electrons. - Electrostatic attraction holds atoms together. - more electron pairs shared, the stronger the bond |
| Octet rule | > atoms are most stable with 8 electrons in their valance shell. |
| Single covalent bond | > ONE electron from each atom is shared so that each now poses a full outer shell. |
| Lone pair electrons | > Pairs of non bonding electrons. |
| Polarity | > when the bond has a slight positive end and a slight negative end, it is polar. - Atoms with similar electronegativities form non-polar covalent bonds. |
| Polar molecules | > asymmetrical molecules. - contains polarized bonds. |
| Non-polar molecules | > symmetrical molecules. - May/may not contain polarized bonds. eg. methane |
| Intramolecular forces | > forces hold atoms of a molecule together. |
| Intermolecular forces | > forces that attract or hold one molecule to another. eg. dispersion forces, dipole-dipole forces, hydrogen bond |
| Dispersion forces | > present between all atoms. - large molecules = large surface area = more dispersion forces = higher boiling point |
| Dipole Dipole forces | > Positive end & negative ends of polar Molecules attract each other. - relatively weak. - more strong polar molecule -- stronger dipole dipole forces. |
| Hydrogen bonds | > are dipole-dipole bonds that ONLY OCCCUR when HYDROGEN is bonded to NITROGEN, OXYGEN, or FLOURINE. - positive H atom -- lone pair NOF - Strongest intermolecular force but weaker than a covalent bond.. |
| Carbon | - tetravalent – it can form 4 covalent bonds with different elements. - allotrope. |
| Diamond | - hardest naturally occurring substance. - consists of carbon atoms bonded very strongly to 4 other carbon atoms in a 3D lattice. - non-conductive. - very high melting point. - insoluble |
| Graphite | - consists of layers of carbon atoms. - layers are held together only by weak dispersion forces. - carbon bonds very strong, between layers not very strong. - can conduct electricity in its solid state. |
| Metallic bonding | > The atoms become positive ions, (cations) because of the delocalized electrons, and become electrostatically attracted to them. - The outer electrons separate from their atoms and become delocalised, creating a ‘sea of electrons’. |
| Melting & Boiling Point (metal) | > metals generally have a very high melting & boiling point. - Metallic bonds are very strong and so a large amount of energy (heat) is needed to break them. |
| Conductivity of Electricity (metal) | > metal generally good conductor of electricity. - the free electrons can carry an electrical charge (moving towards a positive electrode and away from a negative one, creating a circuit.) |
| Conductivity of Heat (metal) | > metal generally good conductor of heat. - The free electrons can take in heat energy, which makes them move faster. They can then transfer the energy throughout the lattice. |
| Malleable and Ductile (metal) | > Metals are usually tough, not brittle. Can be shaped by beating/rolling, can be stretched into wires. - When a metal is hit, the layers of the lattice just slide over each other. The metallic bonds do not break because the electrons are free to move. |
| Lustre (metal) | > metal shiny. - The free electrons in the lattice allow metals to reflect light and appear shiny. |
| Density (metal) | > metals are dense (lol) - The ions in a metal lattice are closely packed. Also depends on the mass of ions, their radius, and how they are arranged. |
| Ionic bonding | > metal & non-metal attracion. |
| Crystal lattice | > many cations & anions bonded together to form = 3D crystal lattice. > electrostatic force between the anions & cations hold ions in place. |
| Ionic Compounds (properties) | > high melting points (the ionic bonds take a lot of energy to break) > hardness & brittleness > electrical conductivity (depends on physical state) |
| Precipitation reaction | > when two ionic and soluble solutions are mixed and a solid forms (precipitate) |
| SNAPE rule (if compound soluble) | > Sodium (Na) > Nitrate (NO₃) > Ammonium (NH₄) > Potassium (K) > Ethanoate (CH₃COO) |
| Relative atomic mass | > mass of each atom compared to a single atom of carbon-12. |
| Relative isotopic mass | > mass of a single atom of that isotope relative/compared to the mass of an atom of carbon-12. |
| Relative molecular mass | > average mass of a molecule of an element/compound relative to the mass of an atom of carbon-12. |
| Mole | > 1 mole of an element is it's mass in grams. - for compounds, you just add the mass of all elements in it together to get what 1 mole of that compound is. |
| Avogadro's constant | > the number of atoms or molecules that are in a mole of a substance. - 6.02 x 10^23 > N = n x Nₐ (N - total number of particles, n - number of mol, Nₐ - Avogadro's constant) |
| Empirical formula | > simplest whole number ratio of th eatoms present in the compound. |
| Molecular formula | >actual number of atoms present in one molecule of the compound. eg. Butane has the molecular formula C₄H₁₀ & empirical formula C₂H₅ |
| Homologous series | > a family of organic compounds eg. alkanes, alkenes... |
| Alkane | > saturated (only single bonds between carbons) > ends in '-ane' formula: CₙH₂ₙ₊₂ |
| Alkenes | > unsaturated (has double bounds between carbons) > ends in '-ene' formula: CₙH₂ₙ |
| Structural isomers | > same molecular formula but a different STRUCTURAL formula. |
| Hydroxyl | > all alcohols contain at least an OH |
| Alcohol | > end in '-ol' > dissolve in water (forms hydrogen bonds) > undergo combustion/burn in air = CO2 & H2O > used as fuels, solvents & drinks |
| Haloalkanes | > dipole-dipole between haloalkanes > somewhat soluble but gets less so with longer carbon chains |
| Carboxyl | > carboxylic acids contain COOH |
| Carboxylic acids | >end in '-oic acid' > dissolve in water = acidic solution > weak acids > more soluble than alcohol > high BP > hygrogen bonds between carboxylic acids |
Created by:
yaannica
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