Some important definitions of Chemistry

Terms	Definitions
Alloy	A compound made from two or more metals
Atom	The tiny particles which make up an element
Atomic number	The number of protons in the nucleus of an atom. All atoms of the same element have the same atomic number
Halogens	Elements in Group 7 of the periodic table.
Group	Term which is applied to the columns on the Periodic Table, also tells you the number of electrons on the outside shell.
Period	Term which is applied to the rows on the periodic table, also tells you the number of shells the atom has.
Compound	Two or more chemical elements joined together.
Conductor	Something that will allow electricity/heat to pass through it.
Diffusion	Net movement of particles by random motion from an area of higher concentration to one of a lower concentration.
Electrons	Negatively charged particles found in atoms.
Elements	Substances that cannot be broken down into simpler substances, for example carbon and oxygen.
Empirical formula	Simplest formula for a substance showing the ratio of elements in it.
Gas	State of a substance where particles move freely and fill the container.
Liquid	State of a substance with no fixed shape but a define volume.
Melt	Change of state from a solid to liquid.
Metal	A shiny substance that conducts electricity.
Neutron	Uncharged particle found in the nucleus of an atom.
Periodic table	Arrangement of elements in order of increasing atomic number, elements with similar properties appearing in the same column.
Product	A substance made in a chemical reaction.
Proton	Positively charged particle found in the nucleus of an atom.
Proton number	The number of protons in the nucleus of an atom. All atoms of the same element have the same atomic number.
Reactant	A substance at the start of a chemical reaction.
Relative atomic mass (Ar)	The mass of an atom compared to other atoms. Ar for hydrogen is 1, and Ar for carbon is 12, so carbon is 12 times heavier than a hydrogen atom.
Relative formula mass (Mr)	The sum of all relative atomic masses of the atoms making up a compound.
Solid	State of matter with fixed volume and shape.
Shell	The position in which electrons orbit a nucleus of an atom.
Transition metal	Elements which sit between Group 2 and 3 of the periodic table.

Notes On Flame Test

Water - Solubility - Ions

FlameTests.

Some metal ions can be identified by the colour of their flame
during a flame test.Metals or metal salts
having a coloured flame are used in fireworks.

A flame test uses a piece of nichrome wire.
You dip the end of the wire inconcentrated hydrochloric acidand then hold it in a hot bunsen flame. If the wire isnot contaminated, the colour of the flamewill not change.
If the flame changes colour, dip thewire in
concentrated hydrochloric acid again and return
it to the bunsenflame.Repeat this procedure
until thewire shows no change of colour in theflame.
Now dip the end of thewire in concentrated hydrochloric acid
and put the wire into the solid which you are
using for the flame test.
A small amount of the metal chloride will form on the wire.
Hold the wire in the flame and see what colour is produced.
The table below shows
which colours are produced by some metal ions.

Metal Ion	Flame Colour
Lithium Li+	Red
Sodium Na+	Yellow/Orange
Potassium K+	Lilac
Calcium Ca2+	Brick Red
Barium Ba2+	Light Green
Copper Cu2+	Blue/Green

flame test, test used in the identification of certain metals. It is based on the observation that light emitted by any element gives a unique spectrum when passed through a spectroscope. When a salt of the metal is introduced into a Bunsen burner flame, the metallic ion produces characteristic color in the flame. Some metals and the colors they produce are: barium, yellow-green; calcium, red-orange; copper salts (except halides), emerald green; copper halides or other copper salts moistened with hydrochloric acid, blue-green; lithium, crimson; potassium, violet; sodium, yellow; and strontium, scarlet. The value of this simple flame test is limited by interferences (e.g., the barium flame masks calcium, lithium, or strontium) and by ambiguities (e.g., rubidium and cesium produce the same color as potassium). A colored glass is sometimes used to filter out light from one metal; for instance, blue cobalt glass filters out the yellow of sodium.

A flame test is an analytic procedure used in chemistry to detect the presence of certain elements, primarily metal ions, based on each element's characteristic emission spectrum. The color of flames in general also depends on temperature; see flame color.

The test involves introducing a sample of the element or compound to a hot, non-luminous flame, and observing the color that results. Samples are usually held on a platinum wire cleaned repeatedly with hydrochloric acid to remove traces of previous analytes.^[1] Different flames should be tried to avoid wrong data due to "contaminated" flames, or occasionally to verify the accuracy of the color. In high-school chemistry courses, wooden splints are sometimes used, mostly because solutions can be dried onto them, and they are inexpensive. Nichrome wire is also sometimes used.^[1] When using a splint, one must be careful to wave the splint through the flame rather than holding it in the flame for extended periods, to avoid setting the splint itself on fire. The use of cotton swab^[2] or melamine foam (eraser)^[3] as a support have also been suggested. Sodium is a common component or contaminant in many compounds and its spectrum tends to dominate over others. The test flame is often viewed through cobalt blue glass to filter out the yellow of sodium and allow for easier viewing of other metal ions.

The flame test is fast and easy to perform, and does not require any equipment not usually found in a chemistry laboratory. However, the range of detected elements is small, and the test relies on the subjective experience of the experimenter rather than any objective measurements. The test has difficulty detecting small concentrations of some elements, while too strong a result may be produced for certain others, which tends to drown out weaker signals.

Although the flame test only gives qualitative information, not quantitative data about the actual proportion of elements in the sample, quantitative data can be obtained by the related techniques of flame photometry or flame emission spectroscopy. Flame Atomic absorption spectroscopy Instruments, made by e.g. PerkinElmer or Shimadzu, can be operated in emission mode according to the instrument manuals.^[4]

Contents 1Common elements 2See also 3References 4External links

Common elements

Some common elements and their corresponding colors are:

Symbol	Name	Color	Image
As	Arsenic	Blue
B	Boron	Bright green
Ba	Barium	Pale/Apple green
Ca	Calcium	Brick red
Cs	Caesium	Blue-Violet
Cu(I)	Copper(I)	Bluish-green
Cu(II)	Copper(II) (non-halide)	Green
Cu(II)	Copper(II) (halide)	Blue-green
Fe	Iron	Gold
In	Indium	Blue
K	Potassium	Lilac
Li	Lithium	Red
Mn (II)	Manganese (II)	Yellowish green
Mo	Molybdenum	Yellowish green
Na	Sodium	Intense yellow
P	Phosphorus	Pale bluish green
Pb	Lead	Blue/White
Ra	Radium	Crimson red
Rb	Rubidium	Red-violet
Sb	Antimony	Pale green
Se	Selenium	Azure blue
Sr	Strontium	Crimson
Te	Tellurium	Pale green
Tl	Thallium	Pure green
Zn	Zinc	Colorless (sometimes reported as bluish-green)

Reactions of acids

You need to be able to describe the reactions of acids with bases, carbonates and metals. You should be able to work out the particular salt formed in the reaction.

Acids and bases

When acids react with bases, a salt and water are made. This reaction is called neutralisation. In general:

acid + metal oxide →salt + water

acid + metal hydroxide→salt + water

Remember that most bases do not dissolve in water. But if a base can dissolve in water, it is also called an alkali.

Carbonates

When acids react with carbonates, such as calcium carbonate (found in chalk, limestone and marble), a salt, water and carbon dioxide are made. In general:

acid + metal carbonate→salt + water + carbon dioxide

Notice that an extra product - carbon dioxide - is made. It causes bubbling during the reaction, and can be detected using limewater. You usually see this reaction if you study the effects of acid rain on rocks and building materials.

Reactive metals

Acids will react with reactive metals, such as magnesium and zinc, to make a salt and hydrogen. In general:

acid + metal →salt + hydrogen

The hydrogen causes bubbling during the reaction, and can be detected using a lighted splint. You usually see this reaction if you study the reactivity series of metals.

Acids, alkalis and neutralisation - higher

When atoms or groups of atoms lose or gain electrons, charged particles called ions are formed. Ions can be either positively or negatively charged.

For the Higher Tier, you need to know which ions are produced by acids, and which are produced by alkalis. You will also need to know the ionic equation for neutralisation.

Acids

When acids dissolve in water they produce hydrogen ions, H⁺. For example, looking at hydrochloric acid:

HCl(aq) → H⁺(aq) + Cl^-(aq)

Remember that (aq) means the substance is in solution.

Alkalis

When alkalis dissolve in water they produce hydroxide ions, OH^-. For example, looking at sodium hydroxide:

NaOH(aq) → Na⁺(aq) + OH^-(aq)

Ammonia is slightly different. This is the equation for ammonia in solution:

NH₃(aq) + H₂O(l) →(aq) + OH^-(aq)

Be careful to write OH^- and not Oh^-.

Neutralisation reaction

When the H⁺ ions from an acid react with the OH^- ions from an alkali, a neutralisation reaction occurs to form water. This is the equation for the reaction:

H⁺(aq) + OH^-(aq) → H₂O(l)

If you look at the equations above for sodium hydroxide and hydrochloric acid, you will see that there are Na⁺ ions and Cl^- ions left over. These form sodium chloride, NaCl.

Salt preparation

You need to be able to work out which particular salt is made in a reaction. You may be asked to describe how to make a salt.

Naming salts

The name of a salt has two parts. The first part comes from the metal in the base or carbonate, or the metal itself if a reactive metal like magnesium or zinc is used.

The second part of the name comes from the acid used to make it. The names of salts made from hydrochloric acid end in -chloride, while the names of salts made from sulfuric acid end in -sulfate.

Formation of salts

Metal		Acid		Salt
sodium hydroxide	reacts with	hydrochloric acid	to make	sodium chloride
copper oxide	reacts with	hydrochloric acid	to make	copper chloride
sodium hydroxide	reacts with	sulfuric acid	to make	sodium sulfate
zinc oxide	reacts with	sulfuric acid	to make	zinc sulfate

Ammonia forms ammonium salts when it reacts with acids. Therefore:

ammonia reacts with hydrochloric acid to make ammonium chloride

Making salts

If the base dissolves in water, you need to add just enough acid to make a neutral solution - check a small sample with universal indicator paper - then evaporate the water. You get larger crystals if you evaporate the water slowly.

Copper oxide, and other transition metal oxides or hydroxides, do not dissolve in water. If the base does not dissolve in water, you need an extra step. You add the base to the acid until no more will dissolve and you have some base left over (called an excess). You filter the mixture to remove the excess base, then evaporate the water in the filtrate to leave the salt behind.

Salt preparation

You need to be able to work out which particular salt is made in a reaction. You may be asked to describe how to make a salt.

Naming salts

The name of a salt has two parts. The first part comes from the metal in the base or carbonate, or the metal itself if a reactive metal like magnesium or zinc is used.

The second part of the name comes from the acid used to make it. The names of salts made from hydrochloric acid end in -chloride, while the names of salts made from sulfuric acid end in -sulfate.

Formation of salts

Metal		Acid		Salt
sodium hydroxide	reacts with	hydrochloric acid	to make	sodium chloride
copper oxide	reacts with	hydrochloric acid	to make	copper chloride
sodium hydroxide	reacts with	sulfuric acid	to make	sodium sulfate
zinc oxide	reacts with	sulfuric acid	to make	zinc sulfate

Ammonia forms ammonium salts when it reacts with acids. Therefore:

ammonia reacts with hydrochloric acid to make ammonium chloride

Making salts

If the base dissolves in water, you need to add just enough acid to make a neutral solution - check a small sample with universal indicator paper - then evaporate the water. You get larger crystals if you evaporate the water slowly.

Copper oxide, and other transition metal oxides or hydroxides, do not dissolve in water. If the base does not dissolve in water, you need an extra step. You add the base to the acid until no more will dissolve and you have some base left over (called an excess). You filter the mixture to remove the excess base, then evaporate the water in the filtrate to leave the salt behind.

8.1 The characteristic properties of acids and bases

1 Describe neutrality and relative acidity and alkalinity in terms of pH (whole numbers only) measured using full-range indicator and litmus.
All substances are divided into three categories:

• Acidic
• Alkaline
• Neutral

How can this be measured?
We usually do this by measuring the pH of the substance. What the pH is that its simply measure of the Hydrogen ion concentration in a substance. However, calculations of that is beyond the scope of the IGCSE Science – if you do, however, want to get a feel of pH calculations, you can visit here.
We measure pH using the pH scale.

• pH 1-6 substances are usually acidic
• pH 7 substances are usually neutral
• pH 7-14 substances are usually alkaline

Universal Indicator

This is a substance that changes color when it is added to another substance. What color it changes to depends on the pH of the substance.

Above: A diagram of a universal indicator

Litmus Paper

This is an indicator also used to test for acidity, neutrality or alkalinity in a substance.

We use something called litmus paper to test for this.

If we want to test for acidity, we use Blue Litmus Paper

If we want to test for alkalinity, we use Red Litmus Paper

The following results are:

• Acids: Turn blue litmus paper red.
• Alkalines/Bases: Turn red litmus paper blue.
• Neutral: No color change.

2 Describe the characteristic reactions between acids and metals, bases (including alkalis) and carbonates.

Metal + Acid → Salt + Hydrogen

We call this the “Displacement” method.
Characteristics of the reaction

• Bubbles are given out
• Temperature rises (the reaction is exothermic, heat is released)
• Metal disappears

Acid + Base → Salt + water

We call this the Neutralization Method. Without fail, water is produced as a product in a neutralization reaction.
There are two types of “Neutralization” reactions.
1) Acid + Metal Oxide → Salt + Water
Copper Oxide + Sulfuric Acid → Copper Sulfate + Water
Here, the Copper merges with Sulfuric acid to make Copper sulfate. If you have iron oxide, nothing will change, the iron will merge with the sulfuric acid to make copper sulfate.
2) Acid + Metal Hydroxide → Salt + Water
Hydrochloric Acid + Sodium Hydroxide → Water + Sodium Chloride
Characteristics of the Reaction
Reaction 1

• Amount of metal oxide decreases
• Temperature increases (exothermic reaction)
• Solution changes color.

Reaction 2

• Hydroxide starts to disappear
• Temperature increases (exothermic reaction)

Acid + Metal Carbonate → Salt + Water + Carbon Dioxide

E.g. Sulfuric Acid (Acid) + Copper Carbonate (Carbonate) → Copper sulfate (salt) + Water + Carbon Dioxide
Characteristics of reaction

• Metal carbonate starts to disappear
• Temperature rises (exothermic reaction)
• Color Change

3 Describe and explain the importance of controlling acidity in the environment (air, water and soil).
Most crops grow best when the pH of the soil is near 7. If soil is too acidic or too alkaline, crops grow badly or not at all.
Usually acidity is the problem. Why? Because of a lot of vegetation rotting in it or because too much fertilizer was used in the past.
To reduce the acidity, the soil is treated with a base like limestone or quicklime or slaked lime.

Affects of lower pH:

• Lack of nutrients
• Poor growth of crops
• May pass onto rivers, damaging the eco-system within it.

Chemical Bonding Chemical compounds are formed by the joining of two or more atoms. A stable compound occurs when the total energy of the combination has lower energy than the separated atoms. The bound state implies a net attractive force between the atoms ... a chemical bond. The two extreme cases of chemical bonds are: Covalent bond: bond in which one or more pairs of electrons are shared by two atoms. Ionic bond: bond in which one or more electrons from one atom are removed and attached to another atom, resulting in positive and negative ions which attract each other. Other types of bonds include metallic bonds and hydrogen bonding. The attractive forces between molecules in a liquid can be characterized as van der Waals bonds. Sodium chloride Ionic Hydrogen molecule Covalent		Index Bond concepts Bond data Chemical concepts
Covalent Bonds Covalent chemical bonds involve the sharing of a pair of valence electrons by two atoms, in contrast to the transfer of electrons in ionic bonds. Such bonds lead to stable molecules if they share electrons in such a way as to create a noble gas configuration for each atom. Hydrogen gas forms the simplest covalent bond in the diatomic hydrogen molecule. The halogens such as chlorine also exist as diatomic gases by forming covalent bonds. The nitrogen and oxygen which makes up the bulk of the atmosphere also exhibits covalent bonding in forming diatomic molecules. Covalent bonding can be visualized with the aid of Lewis diagrams.	Index Bond concepts Chemical concepts
Polar Covalent Bonds Covalent bonds in which the sharing of the electron pair is unequal, with the electrons spending more time around the more nonmetallic atom, are called polar covalent bonds. In such a bond there is a charge separation with one atom being slightly more positive and the other more negative, i.e., the bond will produce a dipole moment. The ability of an atom to attract electrons in the presense of another atom is a measurable property called electronegativity.			Index Bond concepts
Ionic Bonds In chemical bonds, atoms can either transfer or share their valence electrons. In the extreme case where one or more atoms lose electrons and other atoms gain them in order to produce a noble gas electron configuration, the bond is called an ionic bond. Typical of ionic bonds are those in the alkali halides such as sodium chloride, NaCl. Ionic bonding can be visualized with the aid of Lewis diagrams.			Index Bond concepts

Metallic Bonds The properties of metals suggest that their atoms possess strong bonds, yet the ease of conduction of heat and electricity suggest that electrons can move freely in all directions in a metal. The general observations give rise to a picture of "positive ions in a sea of electrons" to describe metallic bonding.

Metal Properties The general properties of metals include malleability and ductility and most are strong and durable. They are good conductors of heat and electricity. Their strength indicates that the atoms are difficult to separate, but malleability and ductility suggest that the atoms are relatively easy to move in various directions. The electrical conductivity suggests that it is easy to move electrons in any direction in these materials. The thermal conductivity also involves the motion of electrons. All of these properties suggest the nature of the metallic bonds between atoms.

Hydrogen Bonding Hydrogen bonding differs from other uses of the word "bond" since it is a force of attraction between a hydrogen atom in one molecule and a small atom of high electronegativity in another molecule. That is, it is an intermolecular force, not an intramolecular force as in the common use of the word bond. When hydrogen atoms are joined in a polar covalent bondwith a small atom of high electronegativity such as O, F or N, the partial positive charge on the hydrogen is highly concentrated because of its small size. If the hydrogen is close to another oxygen, fluorine or nitrogen in another molecule, then there is a force of attraction termed a dipole-dipole interaction. This attraction or "hydrogen bond" can have about 5% to 10% of the strength of a covalent bond. Hydrogen bonding has a very important effect on the properties of water and ice. Hydrogen bonding is also very important in proteins and nucleic acids and therefore in life processes. The "unzipping" of DNA is a breaking of hydrogen bonds which help hold the two strands of the double helix together