chem

HISTORY OF THE ATOM:

Early BCE:

All matter is made of only four fundamental elements: Earth, Fire, Water, and Air. 

Continuum model: Regardless of the number of times you halve a piece of matter, it can always be broken down into smaller pieces. 

460-370 BCE:

Greek philosophers Democritus and Leucippus suggest matter isn’t continuous but made of tiny, solid, unbreakable particles. Clusters of different shapes, arrangements, and positions give rise to macroscopic substances. First use of the term ‘atomos’ meaning indivisible. 

1803: 

English chemist John Dalton developed a more scientific definition of the atom. Small, hard spheres that are indivisible, and that the  atoms of the same element are identical to each other. He theorized that all substances are made of atoms, and atoms are the smallest particles of matter that can’t be divided into smaller particles. Also theorized that atoms of the same elements have the same mass and are alike, different elements have different masses of atoms and are different. 

1897:

British scientist Joseph John Thomson discovered the electron and its negative charge. However, he thought that there must be some positive charge as well in the atom to make the atom charge neutral. 

1904:

The plum pudding model was proposed by Joseph John Thomson (not his term). In this model, an atom is thought of as a round ball of negative charge, with negatively charged electrons embedded in it. 

1911:

New Zealand scientist Ernest Rutherford performed an experiment where he fired a beam of positively charged alpha particles at gold foil.  He found that while most of the alpha particles went through the foil, a small number were deflected.  This led to the development of a nuclear model of the atom in which most of the mass is believed to be contained in a small positive nucleus surrounded by a large space occupied by negative electrons.

1913:

Danish scientist Niels Bohr modified Rutherford’s model proposed that electrons can only travel along certain pathways around the nucleus, called orbits or shells.  This model is sometimes called the planetary model.  This model explained why different elements produce different-coloured light when heated.  This observation is due to the electrons moving from higher to lower orbits and emitting coloured light in the process.

1932:

English scientist James Chadwick discovered the neutron. This showed that the nucleus was not just a mass of positive charge but a cluster of positively charged protons and charge-neutral neutrons.  Further scientific breakthroughs have concluded that the location of an electron in an atom can never be known exactly.  This means that it is impossible for electrons to revolve around the nucleus in specific orbits as suggested by Niels Bohr.  Instead, Erwin Schrodinger suggested that electrons form clouds around the nucleus.  Scientists can predict the shape of these clouds, but never the exact location of electrons within them.

ELEMENTS, COMPOUNDS AND MIXTURES 

Scientists like to classify things and one way to do this is to classify matter by its classification. It's not a good idea to classify by phases (solid, liquid, gas) as it is too vague and phases can change easily. Colour/some other physical aspects aren’t really that defining and are also too vague for items to be classified by them.

Classifying matter questions:

• Is the matter uniform throughout?

• Can it be separated by physical means?

• Can it be separated by chemical means?

Pure substances:

A sample of matter with definite chemical and physical properties. 

Element:

A pure substance that cannot be separated into a simpler substance by physical or chemical means. 

Compound:

A pure substance composed of two or more different elements joined by chemical bonds.

• Made of elements in a specific ratio.

• Has a chemical formula.

• Can only be separated by chemical means, not physical.

Mixture:

A combination of two or more pure substances that are not chemically combined.

• Substances held together by physical forces, not chemical.

• No chemical change takes place.

• Each item retains its properties in the mixture.

• They can be separated physically.

Mixtures can be;

• Well-mixed – Homogeneous

• Not very well-mixed— Heterogeneous

ATOMIC STRUCTURE AND ISOTOPES

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

• Protons (positively charged)

• Neutrons (neutral)

• Electrons (negatively charged)

• Protons and neutrons are relatively equal in 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

• 99.99% of an atom is nothing.

The Rutherford Experiment:

Prior to Rutherford’s experiment, the atom was thought to be a spherical cloud filled with protons and electrons all over (think raisin cakes!)

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

The 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 electrons’ orbits correspond to specific energy levels in the atom.

• Electrons can only occupy fixed energy levels and cannot exist between two energy levels

Scientists quickly extended Bohr’s model of the hydrogen atom to other atoms.

They proposed that electrons were grouped in different energy levels, called electron shells. These electron shells are labelled with the number n = 1, 2, 3 

Different Types of Atoms:

The type of atom that makes 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 

Isotopes:

All atoms that belong to the same element have the same number of protons in the nucleus and therefore the same atomic number, Z. 

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

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

Use of Isotopes:

ELECTRONIC CONFIGURATION:

Using the Bohr Model, it is possible to determine the basic electronic configuration of any atom by applying the following rules:  

Rule 1. Each shell can only contain a maximum number of electrons.

Rule 2. Lower energy shells fill before higher energy shells.

It is important to note that due to the limitations of the Bohr Model, this rule only applies to Elements 1 - 18.

Step 1. Determine the number of electrons (atomic number as for a neutral atom protons = electrons).

Step 2. Recall the maximum number of electrons each shell can hold

Step 3. Place the electrons in the shells from the lowest energy to the highest energy.

Do not exceed the maximum number of electrons allowed!

Step 4. Write the electronic configuration by listing the number of electrons in each shell separated by commas

Valency: In the outer shell (the Valence shell, which contains the Valence electrons), we have to look at the number of electrons that it could lose/gain. For example, a Lithium atom would lose 1 electron due to its electron configuration of 2,1. After this change, we look at the shift in charge for the atom. This is called Valency. This means that Lithium has a valency of +1. 

PERIODIC TABLE INCLUDING PROPERTIES OF GROUPS:

Elements in the periodic table are in order of ATOMIC NUMBER or the number of PROTONS. 

They are divided into;

Periods: Go across, like rows. Same number of occupied electron shells. 

Groups: Go down, like columns. Have the same number of valence electrons. 

Metals are on the LEFT of the table. Nonmetals are on the RIGHT. 

Metal Properties:

• Lustrous

• Good conductors

• High density

• Malleable

• Ductile (can be drawn into wires)

• Usually solid at room temperature

• Opaque as a thin sheet

• Sonorous

Nonmetals Properties:

• Dull

• Poor conductors

• Non ductile

• Brittle

• May be solids, liquids or gases at room temperature

• Transparent as a thin sheet

• Not sonorous

Group 1 - Alkali Metals:

(Hydrogen is not a part of Alkali metals but is part of Group 1)

Valence electron: 1

Chemical Properties: Reacts violently in water that will result in explosions with heavier elements

Physical properties: Soft, shiny and silvery.

Group 2 - Alkaline Earth Metals:

(Harder, less reactive than Group 1. Naturally found in the earth.)

Valence electron: 2

Chemical Properties: Form basic solutions when combined with water

Physical properties: Soft and silvery.

Group 3-12 - Transition Metals:

(Form coloured ions and compounds)

Physical properties: Hard and high density. High melting point and boiling point. 

Group 17 - Halogens:

Reactive non-metals

Valence electrons: 7

Chemical Properties: React with metals to form ionic compounds called salts.

Physical properties: At room temperature fluorine and chlorine are gases, bromine is a liquid. and iodine is a solid

Group 18 - Noble Gases:

Unreactive non-metals due to stable valence shell.

Valence electrons: 8

Chemical Properties: Unreactive (inert). Don’t form ions.

Physical properties: Colourless, odourless and non-flammable.

Metalloids:

Properties of both metals and nonmetals. 

IONS

Ion - a positively or negatively charged atom or group of atoms. When atoms pick up additional electron (s) or lose electron (s), there is no longer a balance between the positive and negative charge. This is an ion.

Ions are formed by the addition of electrons or the removal of electrons from neutral atoms to achieve stability -  i.e a stable valence shell (octet rule). 

• Cations - Positively charged ion
  (atom loses electrons) 
  E.g. Mg²⁺, Al³⁺
  
  

• Anions - Negatively charged ion
  (atom gains electrons)

E.g. Cl^-, O^2-

Neutral charge: An atom that has a neutral charge (i.e. the same number of electrons (negative charge) and protons (positive charge))

Metals have loosely held outer shell (valence) electrons. They can lose these electrons to become cations. This is because the resulting ion has fewer electrons than protons 

CAT-IONS!

Non-metals attract electrons. They gain electrons and acquire a stable outer shell. By gaining electrons, they also increase their negative charge, therefore becoming anions. ANI-ONS!

Naming Ions: 

Cations - Refer to the metal name

E.g. Al³⁺ - Aluminium cation

Anions - Replace the suffix with “-ide”
E.g. O^2- - Oxide anion

Types of bonds:

Remember to follow these three rules:

1. Always write the metal element before the non-metal element in the chemical formula.

2. If the two ions have the same charge (but opposite sign), then they will combine in a one-to-one ratio. The chemical formula will have no subscripts because 1's are never written down. e.g. Na⁺ and Cl^– becomes NaCl, Mg²⁺ and O^2– becomes MgO.

3. If the two ions have different charges, then swap the two numbers and turn them into subscripts. e.g. Ca²⁺ and Cl^– becomes CaCl₂.

Losing/gaining ions requires energy. 

Gain or Loss of Electrons Equations:

E.g. Na -> Na+ + e-

Cl + e- -> Cl-

Mg -> Mg2 + 2e-

O + 2e- -> O2-

ELECTRON TRANSFER AND IONIC BONDING

Ionic Compounds: Usually form when a metal reacts with a nonmetal.

Ionic compound = metallic cation + non-metallic anion

During a reaction there is a transfer of electrons. Metal → Non-metal

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

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.

Transition metals - multiple valencies

Octet Rule: To achieve 8 valence electrons in the outer shell.

Ionic bonds are formed when oppositely charged ions attract (AFTER the electron transfer).

Ionic compounds are unlikely to consist of two positively charged ions as like-charged particles repel.

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(III), Ag(I), Gold (I)
  
  

• The name of the metal cation remains as is
  E.g. Sodium ion, Na⁺ ion
  
  

• Whereas the monatomic, metal anion - substitute the suffix of the name with ‘-ide’