Chemistry 1

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Atoms, elements and compounds

Atoms, elements and compounds

  • All substances are made of atoms. An atom is the smallest part of an element that can exist.
  • Compounds are formed from elements by chemical reactions.
  • Chemical reactions always involve the formation of one or more new substances, and often involve a detectable energy change.
  • Compounds contain two or more elements chemically combined and can be represented by formulae using the symbols of the atoms from which they were formed.
  • Compounds can only be separated into elements by chemical reactions.
  • Chemical reactions can be represented by word equations or equations using symbols and formulae.
  • Name compounds of these elements from given formulae or symbol equations write word equations for the reactions in this specification
  • Write formulae and balanced chemical equations for reactions.
  • Write balanced half equations and ionic equations where appropriate.
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Mixtures

Mixtures

A mixture consists of two or more elements or compounds not chemically combined together.The chemical properties of each substance in the mixture are unchanged.

Mixtures can be separated by physical processes such as filtration, crystallisation, simple distillation, fractional distillation and chromatography. These physical processes do not involve chemical reactions and no new substances are made.

Describe, explain and give examples of the specified processes of separation

Suggest suitable separation and purification techniques for mixtures based on differences in physical properties (eg filtration, crystallisation, distillation, fractional distillation, chromatography).

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The development of the model of the atom

The development of the model of the atom (link with physics)

  • New experimental evidence may change/replace a scientific model.
  • The discovery of the electron led to the plum pudding model of the atom. This suggested that the atom is a ball of positive charge with negative electrons embedded in it.
  • Nuclear model replaced plum pudding model. The mass of an atom is concentrated at the centre (nucleus) and the nucleus is charged.
  • Nuclear model was adapted - electrons orbit the nucleus at specific distances. Scientists Bohr and Chadwick helped to clarify the structure.
  • Experiments showed the positive charge of any nucleus made up of a whole number of smaller particles, each particle (proton) having the same amount of positive charge.
  • New evidence from the scattering experiment led to a change in the atomic model.
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Relative electrical charges of subatomic particles

Relative electrical charges of subatomic particles

Protons     (+1)
Electrons  (-1)
Neutrons   (0)

  • In an atom, the number of electrons = the number of protons in the nucleus.
  • Atoms have no overall electrical charge.
  • The number of protons in an atom of an element = atomic number.
  • All atoms of a particular element have the same number of protons.
  • Atoms of different elements have different numbers of protons.
  • Use the nuclear model to describe atoms.

Atomic number is the smaller number shown in Periodic Table and gives the number of protons (and electrons) in the atom

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Size and mass of atoms

Size and mass of atoms

Atoms are very small, nuclei are even smaller.

Almost all of the mass of an atom is in the nucleus.

Most of the atom is empty space.

Relative masses: protons (1), neutrons (1) and electrons (negligible)

Mass number is the bigger number shown for elements in the Periodic Table (sum of protons + neutrons)

Atoms of the same element can have different numbers of neutrons; these atoms are called isotopes of that element.

Represent examples of atoms, relate size and scale to objects in physical world

 

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Relative atomic mass

Relative atomic mass

The relative atomic mass of an element is an average value that takes account of proportions of the isotopes of the element.

Calculate the relative atomic mass of an element given the percentage abundance of its isotopes.

Example:

Chlorine gas consists of 2 common isotopes with mass numbers 35 (75%) and 37 (25%). The lighter isotope is more abundant.

Average (relative atomic mass) = [(75 x 35) + (25 x 37)]/100

Relative atomic mass of chlorine = 35.3 (usually shown as 35.5 in questions or the Periodic Table)

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Periodic Table

Before atomic structure was understood, known elements were listed in order of atomic weights. Early Periodic Tables had gaps and some elements were in the wrong place because scientists did not then know about isotopes which affected atomic weight. Mendeleev realised he needed to leave gaps for elements as yet undiscovered. He predicted the properties they would have and these elements were later discovered.

In the modern Periodic Table, elements are arranged in order of atomic (proton) number.

Elements with similar properties are found in columns, known as groups. Elements in the same group have the same number of electrons in their outer shell (outer electrons). It is the outer electrons that makes them chemically similar.

Metals are generally found on the left/towards the bottom of the P.T. They react to form positive ions.

Non-metals do not form positive ions and are found towards the top/right hand side.

Be able to explain differences between metals and non-metals and predict possible reactions and probable reactivity of elements from their positions in the P.T.

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Group 0

Group 0

These elements are called the noble gases. They are unreactive and do not easily form compounds because their atoms have stable arrangements of electrons.

Helium He, the lightest noble gas, has 2 electrons filling its outer (only) shell

The other noble gases have 8 electrons in their outer shell.

The boiling points of the noble gases increase going down the group - because weak forces between the atoms increase with size/mass.

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Group 1

Group 1

These elements are known as the alkali metals.

They have characteristic properties because they all only have 1 electron on the outer shell.

They get more reactive going down the group (because the outer electron is more easily transferred during reactions - further away from the positive nucleus)

Be able to describe the reactions of lithium Li, sodium Na and potassium K:

  • with oxygen (metal oxide formed)
  • chlorine(metal chloride formed)
  • water (metal hydroxide and hydrogen formed)

Be able to relate reactivity to 1 electron in outer shell.

Be able to predict properties from trends down the group.

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Group 7

Group 7

These elements are known as the halogens. They exist as simple covalent molecules eg F2, Cl2, Br2, I2 with strong covalent bonds holding the atoms together but weak forces between the molecules.

They have similar rections because they all have 7 electrons in their outer shell.

The further down the group, the higher the melting and boiling points (because higher mass and intermolecular forces become stronger)

They become less reactive going down the group (because the positive nucleus is further away from the incoming electron and so attracts it less strongly).

A more reactive halogen can displace a less reactive halogen from an aqueous solution of its salt.

Example:    Cl2(g)  +  2KBr (aq)  =   Br2(g)   +  2KCl (aq)

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Transition metals

Transition metals

Transition metals are found in long rows towards the bottom of the Periodic Table.They have similar properties to each other, but different from those of Group 1 metals.

Their electron arrangements are more complicated than those of Group 1, but they tend to have more outer electrons and to be able to lose different numbers of those outer electrons.

Compared with Group 1 metals, transition metals:

  • have higher melting points (stronger metallic bonding because of more outer electrons to become delocalised)
  • are stronger and harder (reason as above)
  • have higher densities (heavier atoms)
  • less reactive eg with oxygen, water and halogens (more outer electrons to lose)
  • also form coloured compounds

Be able to show these general properties for chromium Cr, manganese Mn, iron Fe, cobalt Co, nickel Ni and copper Cu.

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The three states of matter

The three states of matter

  • Solid, liquid and gas. Melting and freezing take place at the melting point; boiling and condensing take place at the boiling point.
  • Particle theory helps to explain melting, boiling, freezing and condensing.The amount of energy needed to change state from solid to liquid and from liquid to gas depends on the strength of the forces between the particles of the substance.
  • The stronger the forces between the particles the higher the melting point and boiling point of the substance.
  • Predict the states of substances at different temperatures given appropriate data
  • Explain the different temperatures at which changes of state occur in terms of energy transfers and types of bonding
  • Interpret cooling/heating graphs. Temperature does not change during change of state - eg on melting/boiling, the extra heat energy is being used to separate particles
  • Explain the limitations of the particle theory in relation to changes of state when particles are represented by solid inelastic spheres which have no forces between them.
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State symbols

State symbols

In chemical equations, the three states of matter are shown as (s), (l) and (g), with (aq) for aqueous solutions.

Be able to include appropriate state symbols in chemical equations for reactions

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Electronic structure

Electronic structure

Electrons in an atom occupy the lowest available energy levels (innermost available shells).

The electronic structure of an atom can be represented by numbers or by a diagram. For example, the electronic structure of sodium is 2,8,1

Answer questions in terms of either energy levels or shells.

It is the number of electrons in the outermost shell that determine the reactivity of the element. Electrons closer to the nucleus are attracted more strongly/less easily transferred.

Elements with 1, 2 or 3 outer electrons will be metals (on the left hand side of the Periodic Table)

Elements with 4 or more outer electrons will be non-metals (on the right hand side)

Elements with full outer shells will be in Group 0 (noble gases) and not reactive

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Chemical bonding

Chemical bonding

There are three types of strong chemical bonds: ionic, covalent and metallic.

Ionic bonding - oppositely charged ions attracted strongly together. Occurs in compounds formed from metals combined with non-metals.

Covalent bonding - atoms which share pairs of outer electrons. Occurs in most non-metallic elements and in compounds of non-metals.

Metallic bonding - atoms which share delocalised outer electrons. Occurs in metallic elements and alloys (mixtures of metals).

Be able to explain chemical bonding in terms of electrostatic forces and the transfer or sharing of electrons.

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Ionic bonding

Ionic bonding

When a metal atom reacts with a non-metal atom, electrons in the outer shell are transferred. Metal atoms lose outer electrons to become positively charged ions (with a full outer shell - the electronic structure of a noble gas). The further the outer electrons are away from the nucleus, the more easily lost (metals more reactive going down the group)

Non-metal atoms gain electrons to become negatively charged ions (with a full outer shell). The closer the electrons are to the nucleus, the greater the attraction (non-metals more reactive going up the group).

Electron transfer can be represented by dot and cross diagrams (be able to draw these for ionic compounds formed by metals in Groups 1 & 2 with non-metals in Groups 6 & 7).

Be able to work out the charge on the ions of metals and non-metals from the group number of the element (metals in Groups 1 and 2, and non-metals in Groups 6 and 7).

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Ionic compounds

Ionic compounds

An ionic compound is a giant structure of ions. Ionic compounds are held together by strong electrostatic forces of attraction between oppositely charged ions. These forces act in all directions in the lattice.

Be able to:

  • deduce that a compound is ionic from a diagram of its structure
  • describe the limitations of using dot and cross, ball and stick, two and three-dimensional diagrams to represent a giant ionic structure
  • work out the empirical formula of an ionic compound from a given model or diagram that shows the ions in the structure.
  • be familiar with the structure of sodium chloride (do not need to know the structures of other ionic compounds).
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Properties of ionic compounds

Properties of ionic compounds (eg sodium chloride NaCl)

Ionic compounds have regular structures (giant ionic lattices) in which there are strong electrostatic forces of attraction in all directions between oppositely charged ions (eg Na+ and Cl-).

These compounds have high melting points and high boiling points because of the large amounts of energy needed to break these strong bonds.

When melted or dissolved in water, ionic compounds conduct electricity because the ions are free to move and so charge can flow.

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Covalent bonding

Covalent bonding

When atoms share pairs of electrons, they form strong covalent bonds.

  • Covalently bonded substances may consist of small molecules (eg O2).
  • Recognise common substances consisting of small molecules from chemical formula.
  • Draw dot and cross diagrams for the molecules of hydrogen, chlorine, oxygen, nitrogen, hydrogen chloride, water, ammonia and methane
  • Some covalently bonded substances have very large molecules, such as polymers.
  • Some covalently bonded substances have giant covalent structures, such as diamond, graphite and silicon dioxide.
  • Represent the covalent bonds in small molecules, in the repeating units of polymers and in part of giant covalent structures, using a line to represent a single bond
  • Describe the limitations of using dot and cross, ball and stick, two and three-dimensional diagrams to represent molecules or giant structures
  • Deduce the molecular formula of a substance from a given model or diagram in these forms showing the atoms and bonds in the molecule.
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Properties of small molecules

Properties of small molecules

'Molecule' always means atoms are bonded together with strong covalent bonds

Substances consisting of small molecules are usually gases or liquids with relatively low melting points and boiling points (eg H2, Cl2, O2, N2, HCl, H2O, NH3, CH4)

They have only weak forces between the molecules (intermolecular forces). It is these intermolecular forces that are overcome, not the covalent bonds, when the substance melts or boils.

The intermolecular forces increase with the size of the molecules, so larger molecules have higher melting and boiling points.

Covalent substances do not normally conduct electricity because the molecules do not have an overall electric charge.

Use the idea that intermolecular forces are weak compared with strong covalent bonds to explain the properties of molecular substances.

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Polymers

Polymers

Examples are polythene (polyethene), polystyrene, nylon etc

Polymers have very large molecules.

The atoms in the polymer molecules are linked to other atoms by strong covalent bonds.

The intermolecular forces between polymer molecules are relatively strong (because they are very large molecules) and so these substances are solids at room temperature.

Be able to recognise polymers from diagrams showing their bonding and structure.

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Giant covalent structures

Giant covalent structures

These are solids with very high melting points. All of the atoms in these structures are linked to other atoms by strong covalent bonds. These bonds must be overcome to melt or boil these substances. Diamond and graphite (forms of carbon) and silicon dioxide (silica) are examples of giant covalent structures.

Recognise giant covalent structures from diagrams showing bonding and structure.

Diamond is a giant lattice where all the carbon atoms use all their outer electrons to form 4 strong covalent bonds with other carbon atoms (in a tetrahedral arrangement). High mp, non-conductor, extremely hard (but brittle and can be cleaved if hit with a sharp metal edge).

In graphite, each carbon atom forms 3 covalent bonds with three other carbon atoms, forming layers of strongly bonded hexagonal rings but weaker intermolecular bonding between the layers.The layers can be easily separated (eg in pencil 'lead' and as a lubricant). One electron from each carbon atom is delocalised and can move between the layers. Therefore graphite can conduct electricity (and heat) in the direction parallel to the layers. Graphite is similar to metals in that it has delocalised electrons.

 

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Graphene and fullerenes

Graphene and fullerenes are exciting new giant covalent forms of carbon.

Graphene is a single layer of graphite and has properties that make it useful in electronics and composites (flexible but very strong, conducts electricity). Explain the properties of graphene in terms of its structure and bonding.

Fullerenes are huge molecules of carbon atoms with hollow shapes/cages. The structure of fullerenes is based on hexagonal rings of carbon atoms but they may also contain rings with five or seven carbon atoms (like a football). The first fullerene to be discovered was Buckminsterfullerene (C60) which has a spherical shape.

Carbon nanotubes are very long, thin cylindrical fullerenes. Their properties make them useful for nanotechnology, electronics and materials.

  • Recognise graphene and fullerenes from diagrams and descriptions of their bonding and structure
  • Give examples of the uses of fullerenes, including carbon nanotubes eg other molecules can fit into the hollow areas inside fullerenes eg medicines to be transported to a particular organ in the body.
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Metallic bonding

Metallic bonding

  • Metals consist of giant structures of atoms arranged in a regular pattern (lattice).
  • The electrons in the outer shell of metal atoms are delocalised and so are free to move through the whole structure. The sharing of delocalised electrons gives rise to strong metallic bonds.
  • Be able to explain how bonding and structure are related to the properties of metals
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Properties of metals and alloys

Properties of metals and alloys

Metals have giant structures of atoms with strong metallic bonding.

This means that most metals have high melting and boiling points.

In pure metals, atoms are arranged in regular layers, which allows metals to be bent/shaped (malleable). Pure metals are too soft for many uses and so are mixed with other metals to make alloys which are harder (eg pure gold is quite soft, mixed with silver, copper, platinum to harden for jewellery)

Be able to explain why alloys are harder than pure metals in terms of distortion of the layers of atoms when atoms of different sizes are mixed in.

Metals are good conductors of electricity because the delocalised electrons (the outer electrons of each atom) in the metal carry electrical charge through the metal.

Metals are good conductors of thermal energy because energy is transferred by the delocalised electrons.

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Sizes of particles and their properties

Sizes of particles and their properties

Nanoparticles are structures that are 1–100 nm in size. (1nm = 10-7 m)

Fine particles have diameters between 100 and 2500 nm.

Coarse particles have diameters between 1 x 10-5 m and 2.5 x 10-6 m. Coarse particles are often referred to as dust.

As the side of cube decreases by a factor of 10 the surface area to volume ratio increases by a factor of 10.

Nanoparticles may have properties different from those for the same materials in bulk because of their high surface area to volume ratio. It may also mean that smaller quantities are needed to be effective than for materials with normal particle sizes (NB larger SA increases rates of reaction).

Students should be able to compare ‘nano’ dimensions to typical dimensions of atoms and molecules.

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Uses of nanoparticles

Uses of nanoparticles

Nanoparticles have many applications in medicine, in electronics, in cosmetics and sun creams, as deodorants, and as catalysts. New applications for nanoparticulate materials are an important area of research.

Consider advantages and disadvantages of the applications of these nanoparticulate materials, but do not need to know specific examples or properties other than those specified.

Be able to:

•given appropriate information, evaluate the use of nanoparticles for a specified purpose

•explain that there are possible risks associated with the use of nanoparticles.

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Conservation of mass and balanced chemical equatio

Conservation of mass and balanced chemical equations

The law of conservation of mass states that no atoms are lost or made during a chemical reaction so the mass of the products must equal the mass of the reactants.

This means that chemical reactions can be represented by symbol equations, balanced in terms of the numbers of atoms of each element involved on both sides of the equation.

Understand the use of the multipliers in equations in normal script before a formula and in subscript within a formula.

Relative formula mass

  • The relative formula mass (Mr) of a compound is the sum of the relative atomic masses of the atoms in the numbers shown in the formula (masses will be provided in the exam)
  • In a balanced chemical equation, the sum of all the relative formula masses of the reactants in the quantities shown equals the sum of the relative formula masses of the products in the quantities shown.
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Mass changes when a reactant or product is a gas

Mass changes when a reactant or product is a gas

Some reactions may appear to involve a change in mass but this can usually be explained because a reactant or product is a gas and its mass has not been taken into account.

For example:

  • when a metal reacts with oxygen, the mass of the oxide produced is greater than the mass of the metal
  • in thermal decompositions of metal carbonates, carbon dioxide is produced and escapes into the atmosphere leaving the metal oxide as the only solid product.

Explain any observed changes in mass in non-enclosed systems during a chemical reaction given the balanced symbol equation for the reaction and explain these changes in terms of the particle model.

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Chemical measurements

Chemical measurements

Whenever a measurement is made there is always some uncertainty about the result obtained (sometimes shown as +/-)

Be able to:

  • represent the distribution of results and make estimations of uncertainty
  • use the range of a set of measurements about the mean as a measure of uncertainty
  • interpret graphs, identify anomalous results and predict missing values

 

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Moles

Moles

Chemical amounts are measured in moles. The symbol for the unit mole is mol.

The mass of one mole of a substance in grams is numerically equal to its relative formula mass.

One mole of a substance contains the same number of the stated particles, atoms, molecules or ions as one mole of any other substance.

The number of atoms, molecules or ions in a mole of a given substance is the Avogadro constant. The value of the Avogadro constant is 6.02 x 1023 per mole.

Amounts in moles can apply to atoms, molecules, ions, electrons, formulae and equations, for example that in one mole of carbon (C) the number of atoms is the same as the number of molecules in one mole of carbon dioxide (CO2).

Be able to use the relative formula mass of a substance to calculate the number of moles in a given mass of that substance and vice versa.

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Amounts of substances in equations

Amounts of substances in equations (HT only)

The masses of reactants and products can be calculated from balanced symbol equations.

Chemical equations can be interpreted in terms of moles.

For example:     Mg + 2HCl MgCl2 + H2

shows that 1 mole of magnesium reacts with 2 moles of hydrochloric acid to produce 1 mole of magnesium chloride and  1 mole of hydrogen gas.

Be able to:

  • calculate the masses of substances shown in a balanced symbol equation
  • calculate the masses of reactants and products from the balanced symbol equation and the mass of a given reactant or product.
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Calculating masses in a reaction using moles

Calculating masses in a reaction using moles

Starting from a balanced equation, the masses of reacting substances used or products made can be calculated using moles.

Example:  

Calculate the maximum mass of iron III chloride (FeCl3) that can be made from 11.2g of iron

Balanced equation:       2Fe  +  3Cl2    2FeCl3

                                   2 mol. iron    →   2 mol. iron III chloride
                                   2 x Mr iron    →   2 mol. Mr iron III chloride
                                   2 x 56g         →   2 x (56 + 3x 35.5)g
                                     112g           →    2 x 162.5g (= 325g)

(divide both sides by 10)  11.2g iron would make 32.5g iron III chloride

(You can also work out a balanced equation if you know all the masses)

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Percentage yield

Percentage yield

Even though no atoms are gained or lost in a chemical reaction, it is not always possible to obtain the calculated amount of a product because:

  • the reaction may not go to completion because it is reversible
  • some of the product may be lost when it is separated from the reaction mixture
  • some of the reactants may react in ways different to the expected reaction.

The amount of a product obtained is known as the yield. When compared with the maximum theoretical amount as a percentage, it is called the percentage yield.

% Yield = (Mass of product actually made ÷ Max theoretical mass of product)  x 100              

Be able to calculate the percentage yield of a product from the actual yield of a reaction

Be able to calculate the theoretical mass of a product from a given mass of reactant and the balanced equation for the reaction.

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Atom economy

Atom economy (HT)

The atom economy of a chemical reaction is a measure of the amount of starting materials that become useful products. Inefficient, wasteful processes have low atom economies. Efficient processes have high atom economies, and are important for sustainable development, as they use fewer natural resources and create less waste.

atom economy =  (mass desired product ÷ total mass products) x 100

Example: What is the atom economy of the process to produce hydrogen from coal and steam?

Balanced equation                C(s) + 2H2(g)   CO2(g) + 2H2(g)
Mr of products                                                44g        2 x 2g
Total mass products                                  =    44 + 4 (= 48)g
Atom economy                                         =    (4÷ 48) x 100
                                                               =     8.3%
This is a low atom economy, so an inefficient process (and uses coal, a non-renewable resource)

 

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Limiting reactants

 Limiting reactants

In a chemical reaction involving two reactants, it is common to use an excess of one of the reactants to ensure that all of the other reactant is used. The reactant that is completely used up is called the limiting reactant because it limits the amount of products.

Students should be able to explain the effect of a limiting quantity of a reactant on the amount of products it is possible to obtain in terms of amounts in moles or masses in grams.

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