term 2: chemistry - school notes

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matter and atoms

matter

• a substance that has volume (occupies space) and has a mass

• all matter are made up of atoms which bond together to produce different substances

• matter and energy are not the same thing

• everything is either matter or energy

• energy does not have any atoms

• energy and matter can change either and vice versa

• first element to be found is helium - in the world

atoms

• atom : building block of matter

• 1802: first atomic theory of matter was presented by John Dalton. Dalton proposed that all matter is made up of tiny spherical particles, which are indivisible and indestructible

• we now know it is INCORRECT and atoms are in fact made of smaller subatomic particles (protons, neutrons and electrons)

Pure substances VS Mixtures

• mater can be split into two groups: pure substances and mixtures

Mixtures

• Mixtures: consist of 2 or more types of particles that are not chemically combined together. Can be made elements, compoounds or both

  • air: oxygen gas (O2), Carbon dioxide gas (CO2), Nitrogen gas (N2)

  • Oil-water mixture

• Mixtures can be seperated by physical means (e.g. filtration)

Pure Substances

• Pure substances: A pure substance is a substance made up of only one type of particle throughout. this means its has fixed composition and consistent properties. This can either be one single element or one single compound, but every sample of this substance that you examine must contain exactly the same thing with fixed, definite set of properties

  • pure element: copper metal (cu)

  • Pure compound: carbon dioxide gas (CO2)

• It cannot be seperated by physical means (e.g. filtration)

Elements

• Elements: Made of just one type of atoms

• It could be monatomic (i.e exist as individual atoms) or could also form molecules

• molecules: 2 or more atoms that are held together by chemical bonds

• elements cannot be seperated into simpler substance by physical or chemical means

Compound

• Compound: different type of atoms could combine to form new substances

• To determine whether a pure substance is an element or a compound, you must determine if the substance can be broken down into simpler substances

• Molecules: 2 or more atoms that are covalently bonded

• MOLECULES MUST BE COVALENTLY BONDED

• all molecules are compounds but not all compounds are molecules

• compounds are ionically bonded

Atomic structure

• an atom is made up of three types of sub-atomic particles:

  • protons (positively charged)

  • Neutrons (neutral)

  • Electrons (negatively charged)

• in chemistry, the word particle is a general term that refers to a small unit of matter

• Depending on the context, “particle“ could mean an atom, a molecule, an ion, or something else.

Atomic strctures

• protons and neutrons are relatively equal mass and are huge in comparisions to electrons. This is why the nucleus is so dense

• proton and neutron contribute nearly all the mass of the atom

• Protons and neutrons are relatively equally mass and are huge in comparison to electrons. this is why the nucleus is dense

• Proton and neutron contribute nearly all the mass of the atom

The Rutherford Experiment

• Rutherford’s model proposed the following

  • most of the mass of an atom and all of the positive charge must be located in a tiny central region called the nucleus

  • Most of the volume of an atom is empty space, occupied only by electrons

  • The electrons move in a circular orbits around the nucleus

  • The force of the attraction between the positive nucleus and the negative electrons is electrostatic.

• Prior to Rutherford’s experiment, the atom was thought to be a spherical cloud filled with protons and electrons all over (think raising caked, or plum pudding)

• In 1911, Rutherford’s experiment proved that the atom has a tiny but heavy nucleus and that the most of the volume of an atoms is an empty space occupied by electrons

Bohr model

• In 1913 Niels Bohr developed a new model of the hydrogen atom that explained emission spectra. The Bohr model proposed the following

  • electrons revolve around the nucleus in fixed, circular orbits

  • the electron’s orbits correspond to specific energy levels in the atom

  • Electrons can only occupy fixed energy levels and

• scientists quickly extended Bohr’s model of the hydrogen atom to the other atom

• they proprosed that electrons were grouped in tdiffernt energy levels, called electron shells

• these electron shells are labelled with the number n = 1, 2, 4

  The electron shells

• the highest energy level found is n7

• highest energy level is found farther away from the nucleus because of the electrostatic attraction force

• the higher up you go the higher the electrostatic attraction therefore lower energy the closer you get to the nucleus

• if 2 negatively charged electrons share an electron shell have anti clockwise spins so they never meet or never come in contact with one another

• think 2 siblings circling each other

Different types of atoms

• the type of atom that make up each element is determined by the number of protons (atomic number) in the nucleus

  • atomic number: the number of protons in the nucleus of the atom

  • mass number: the total number of protons + neutrons in the nucleus

  • Atoms are electrically neutral: therefore number of electrons = number of protons

• the number of neutrons is calculated by: atomic mass - atomic number (number of electrons/protons)

• the atomic number is usually the smaller number and the atomic mass is usually larger and has a lot of decimals

Isotopes

• all atoms that belong to the same element have the same number of protons in the nucleus and therefore the same atomic number,

• atoms that have the same number of protons (atomic number) but different numbers of neutrons (and therefore different mass ambers) are known as isotopes

• isotopes have identical chemical properties but different physical properties such as mass and density. in particular, some isotopes are radioactive

• In nature, different elements have different numbers of isotopes. Gold only has one isotope, whereas lead has four isotopes and mercury seven isotopes

• Protons: determine elemental identity

• Electrons: determine chemical reactivity(how easily (or not easily) it is for the atom to undergo change in a chemical reaction )

• Neutrons: determine physical properties the physical properties of an element

Practice with Isotopes

• Identify the most abundant form of carbon out of the 3

1. ¹²₆ C - Carbon 12

2. ¹³₆ C - Carbon 13

3. ¹⁴₆ C - Carbon 14

• Carbon 12 - it appears in the periodic table meaning it is the most commonly found Isotope

• note that 12, 13, 14 are the atomic mass of the atom

Electron configuration

• Using the bohr model it is possible to determine the basic electronic configuration of an atoms by applying the folowing rules

1. eaach shell can only contain a maximum number of electrons

2. lower enegy shells full before hugher ernergy shells

Steps to annotating an atom - Electron configuration

1. determine the number of electrons in your element

2. recall the maximum number of electrons each shell can hold

3. place the number of electrons in the shells from the lowest energy to the highest energy. do no exceed the max number of electrons allowed

4. write the electronic configuration by listing the number of electrons in each shell separated by commas. (2, 8, 1) (for sodium)

• THIS ONLY APPLIES TO THE FIRST 18 ELEMENTS (TILL ARGON)

Atoms

• Atoms: Building block of matter

Ions

• A positively or negatively charged atom or group of atom

• remember than an atom has the same number of electrons and protons which means that is it neutral

• When atoms pick up an additional electron (s) or lose electrons (s), there is no longer a balance between the positive and negative charge

• They lose electrons/protons to achieve a full valence shell (octet rule)

  • an exception to the octet rule would be the duet rule (applies to hydrogen & helium)

  • the duet rules states that there would 2 electrons in the valence shell

2 main types of ions

1. monoatomic e.g.

   a. Na⁺, Cl^-, Cu⁺²

2. Polyatomic

   a. e.g. Pos₄^3-, OH^-

Cations

• a positively charged ion (atom loses electrons)

  • Meg2+, Al3+

• Magnesium for example has 2,8,2

• it loses 2 electrons and becomes

  • Mg ^+2/2+

  • You would also put square brackets around the visual representation and put a 2+/+2 in the upper right corner

Anions

• Negatively chaged ion (atom gains electrons)

  • e.g. Cl-, O2-

• Chlorine for example has 2, 8, 7

• It gains an electrons and becomes

  • Cl^-

  • there is an implied one

Metals

• Metals have loosely held outershells (valence electrons)

• they can lose these electrons to become cations, this is because the resulting ion has fewe electrons than protons

Nonmetals

• non metals are attract electrons

Transition metals

• they can take on more than 1 charge

• they have variable charges

• Really big valence shells so they can lose or gain electrons really easily

positive control : variable with known result (e.g. plant is exposed to sunlight)

negative control: absence of IV (e.g. plant exposed to no light)

• No result (e.g. no change in plant height/mass)

Electron Transfer and Ionic Bonding

Recall classifying matter

atom : smallest particle of matter

Element: a substance made from only type of atom

Compound: A substance made from 2 or more different types of atoms

Molecules: A substance made from two or more atoms that chemically combined

Ionic compound

• An ionic compound typically froms when a metal reacts with a nonmetal

• Ionic compound = metallic cation + non metalli anion

• During a reaction there is a transfer of electrons Metals → Non-metal

• Once this occurs, the oppositely charged ions join together in a lattice

• oppositely charged ions are attracted to one another and this attraction would be called an “electrostatic attraction”

• electrostatic attraction AKA ionic bonds AKA the attraction between oppositely charged ions

• They are brittle as when a force is applied to the lattice the ions with like charges align and actively repel each other shattering the lattice

• The ionic bond is very strong, and brittle because rearranging ions causes repulsion between like charges

• You are sliding the layers so the bond is still intact but the layers slide and the negatives end up next to negatives so they repel each other and then the lattice breaks apart

• they normally have very high mpts

• energy must be transferred to a substance to make it melt or boil

• this energy overcome the strong electrostatic forces of attraction which act tin all directions between the oppositely charged ions

  •  Some forces are overcome during melting

  • all remaining forces are overcome during boiling

• the more energy needed the higher the melting point of boiling points

• since the electrostatic forces of attraction between oppositely charged ions are strong, their melting and boiling points are high

• solid sodium chloride needs a temperature of 801 degrees to melt it

Electron transfer diagrams

• Electron transfer diagrams are used to show the path that electrons take when they are removed from a metal and added to a non-metal during ionic bonding

• show the transfer of electrons from metals to non-metals

Steps to draw an electron transfer diagram

1. Draw an electron shell diagram of the neutral metal and non-metal

2. Add a ‘+‘ between them

3. Draw an arrow leading from each valence in the metal to the valence shell of the non-metal

4. Add an arrow towards the resulting ions

5. Draw an electron shell diagram of the resulting cation/s and anion/s

6. Write the chemical symbol of the metal and the non-metal in the centre of the electron shell diagram, taking care of the electron shell diagram, taking care to add charges to your ions.

Naming Ionic Compounds - rules

• Name cations first (metals) before anions

• If metal is a transition metal, indicate the valency in numerals after the name

  • e.g: Fe (lll), Ag (l), Gold (l)

• the name of the metal cation remains as is

  • e.g: Sodium ion, Na⁺ ion

• the name of the non-metal anion is changed. it’s suffix becomes ‘-ide‘

Shortcut method - e.g.

Al³⁺ + O^2-

Al₂ O₃

Metallic Ions

• can only be made of cations

• metallic bonding occurs in pure metals

• metals occur as crystal lattices

• this is because metallic atoms tend to lose their outer shell electrons easily

• this turns the metals atoms in positively-charged cations

• the cations form a metallic lattice structure, in which electrons from each metal atom overlap with each other, forming a sea of electrons that can flow between all the metal ions

• crystal lattices are very structured and amongst all of that are these free flowing negative charges

• if you draw them draw atoms touching each other

• electrostatic forces of attraction between the positively charger metal cations and negatively charged valence electrons occur in all direction, holding the lattice together

• this type of no-directional bonding is known as metallic bonding

• this means that metal atoms are hard to separate but relatively east to move (malleable)

Alloys

• alloys are mixture of 2 or more elements where 1 elements is a metal, combined via metallic bonding

• alloys are generally harder than the pure elements they contain, this is due to the pure metal atoms being the same and arranged in layers, as opposed to alloys which contain elements with atoms of different sizes

• you do not create compounds when you stick it with 2 metals you only get an alloy

• e.g: steel (iron and carbon), bronze (copper and tin) and brass (copper and zinc)

Types of bondings

Looking closer at ionic susbstance

a stick and ball representation of the lattice structure of sodium chloride ; the sticks represent the bonds.

ionic compounds are solid at room temperature

• inioc compounds are called ‘salts‘ and are solid at room temperature

• ionic compounds have strong electrostatic forces of attraction between their positive and negative ions and form a lattice of ions

• a high amount of energy is required to over come these forces

• therefore, they exist as solids at room temperature

• to shatter an ionic bond requires a high amount of energy

• so inonic bonds are the strongest bons

high melting points

• high melting points in the ionic bonding model result from the strong electrostatic force by oppositely charged ions

• to metl an ionic copimtnd, a large amount fo energy is needed to break this strong bond and allow ions to move freely

• as mentioned before, this means they are solid at room temperaatre

why are ionic compounds brittle

• ionic ompounds are brittle because, when external force is applied to them, the charged ions (cations + ions) shift therefore align with ions

• this causes them to repel eachother, cuasing them to shatter

ionic compounds cna be soluble in water

• water can break aparat the ionic lattice because the positive cations and negative anions are attracted to the partial charges within the water molecules

• the kinetic energy of the water molecules breaks the elctrostatics attractions inside the ionic lattice strucutre

• this allows water to surround the individual ions furing dissolving

conduction of electricity in an aqueous solution

• substances that conduct electricity have free moving, charged particles

• when the charged ions are dissolved in water they can move to conduct electricity

• if you are able to push particles easily through matter you can conduct electricity through it

they are conductive when in molten form

• molten ionic compounds can conduct electricity because they have free mobile charged particles

Accuracy

reproducibility : different conditions but same experiment

repeatability: same conditions in the same experiment

•    goal is to get the same results for both

covalent bonding

• covalent bonding occurs between two nonmetals

• however, atoms still want to achieve a noble gas configuration (the octet rule)

• But, rather than losing or gaining electrons, atoms now share an electron pair

• the shared electron pair is called a bonding pair

• e.g: Cl₂ forms a covalent bond with itself

• Covalent bonds with the same elements are known as diatomic atoms

• theres only one pair of shared electrons

• all halogens can be covalently boded with themselves