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Mg(s) + Zn²⁺ →→ Zn(s) + Mg²⁺
The interval between metals in the reactivity series represents the reactivity of those metals towards each other.
If the interval between elements is larger, they will react more vigorously. The topmost five elements, form lithium to sodium are known as very active metals; hence they react with cold water to produce the hydroxide and hydrogen gas. For example, sodium forms sodium hydroxide and hydrogen gas with cold water.
2Na + 2H₂O →→ 2NaOH + H₂
From magnesium to chromium, elements are considered as active metals and they will react with very hot water or steam and form the oxide and hydrogen gas. For example, aluminum reacts with steam to form aluminum oxide and hydrogen gas.
2Al + 3H₂O →→ Al₂O₃ + 3H₂

From iron to lead, metals can replace hydrogen from various acids like Hydrochloric acid, dilute sulfuric and nitric acids. Oxides of these metals undergo reduction when heated with hydrogen gas, carbon, or carbon monoxide. Till copper, metals can combine directly with oxygen and form metal oxide. Elements present at the bottom from mercury to gold are often found in the native form in nature and their oxides show thermal decomposition under mild conditions.

Reactivity Series Chart
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We can summarize the reactivity of different metals in a reactivity series chart.

Metal	Symbol	Reactivity	Extraction
Lithium	Li
Potassium	K	displaces H₂ gas from water, steam and
Strontium	Sr	displaces H₂ gas from water, steam and
Strontium	Sr	acids and forms hydroxides	Electrolysis
Calcium	Ca	acids and forms hydroxides
Calcium	Ca
Sodium	Na
Magnesium	Mg	displaces H₂ gas from steam and acids
		and forms hydroxides
Aluminium	Al	and forms hydroxides
Carbon	C	Included for comparison
Manganese	Mn	displaces H₂ gas from steam and acids
Zinc	Zn	displaces H₂ gas from steam and acids
Zinc	Zn	and forms hydroxides
Chromium	Cr	and forms hydroxides
Chromium	Cr
Iron	Fe		Smelting with coke
Cadmium	Cd		Smelting with coke
Cobalt	Co	displaces H₂ gas from acids only and
Nickel	Ni	forms hydroxides
Tin	Sn
Lead	Pb
Hydrogen	H2	included for comparison
gas	H2	included for comparison
gas
Antimony	Sb
Arsenic	Ar	combines with O₂ to form oxides and
		cannot displace H₂
Bismuth	Bi	cannot displace H₂
Copper	Cu		Heat or physical
Mercury	Hg		Heat or physical
Mercury	Hg		extraction methods
Silver	Ag	found free in nature, oxides decompose	extraction methods
Silver	Ag
Paladium	Pd
Paladium	Pd	with heating
Platinum	Pt	with heating
Platinum	Pt
Gold	Au

All metals have a tendency to lose electrons and form metal ions. In other

words, all metals are good reducing agents and easily oxidise themselves. M →→ Mⁿ⁺ + ne^-

The reactivity series of elements can be shown in another way, which includes oxidation reaction of each metal to the respective metal ion. It gives information regarding the reducing power of the metal atom and the oxidation number of the

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metal ion.
Exothermic and endothermic reactions
All processes can be classified into one of two categories: exothermic and endothermic. In an exothermic process, energy is released, while in an endothermic process, energy is stored. This section will specifically cover exothermic and endothermic chemical reactions, but almost any process can be described as releasing or storing energy.
The concept of giving off or storing energy can sometimes be a bit confusing, so let's go over some of the basic types of energy that you'll encounter in your chemistry class, and what it means to give off and store each type of energy.
Heat: Heat energy is the energy that accompanies temperature changes. If heat energy is being released then the reaction from which it is released will become hotter. If heat energy is being stored, then reaction will become colder.
Light: If light energy is being given off, then the reaction will glow. If it's being stored, then the reaction will seemingly proceed on its own without any catalyst present without any heat being evolved or absorbed.
Mechanical energy: If mechanical energy is being stored, then the volume and/or pressure of the reaction will get smaller. If mechanical energy is being given off, then the opposite will be true.
The most common change in energy that you'll witness in your chemistry class will be changes in heat energy. It can be measured with a bomb calorimeter. Energy released or stored in a reaction will often be expressed written as H, or a change in enthalpy. A positive H means that energy is stored and the reaction is endothermic. A negative H means that energy is released and the reaction is exothermic. It is usually expressed in kilojoules (kJ) or joules (J).
Why Exothermic Or Endothermic?
If you understand the above section, then you can now identify whether a reaction is exothermic or endothermic. If it gives off one of the above three types of energy then it's exothermic, if it absorbs it, then it's endothermic. The question that still hasn't been answered, though is why? Why are some reactions exothermic while others are endothermic, and why does energy have to be absorbed or released at all?
The answer lies in chemical bonds. Chemical bonds have bond energies associated with them. This bond energy is the amount energy that it takes to break the bonds, and also the amount of the energy that is released when the bonds are formed. Consequently, if the bonds in your reactants have a higher total bond energy than your products, the reaction will be endothermic. If they have a lower total bond energy, it will be exothermic.

The reason for this is the law of conservation of energy, which states that energy cannot be created or destroyed; it can only change forms. In this case, it would mean that whatever energy was used to break the bond will be released if the bond is reformed. For example:

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Suppose we have a C-H bond somewhere, and we wanted to break that bond apart into a C and an H. We'd have to put in some amount of energy. Let's call this amount 'x'. Once we put in x energy, by say, adding heat, the C-H bond will break apart. What happened to 'x' though? The conservation of energy law says that 'x' didn't just disappear; it just took on another form, in this case exciting the electrons in C and H. Some of the energy went to the C atom and some went to the H atom. If the C-H bond reformed, then 'x' would be released again. If the C went off and recombined with a different molecule (let's say a Cl), and so did the H (with an F, for instance). Then the energy released from the new pairings would be 'x' plus whatever energy the Cl and F had stored.
Many chemical reactions release energy in the form of heat, light, or sound. These are exothermic reactions. Exothermic reactions may occur spontaneously and result in higher randomness or entropy ( S > 0) of the system. They are denoted by a negative heat flow (heat is lost to the surroundings) and decrease in enthalpy ( H < 0). In the lab, exothermic reactions produce heat or may even be explosive.
There are other chemical reactions that must absorb energy in order to proceed. These are endothermic reactions. Endothermic reactions cannot occur spontaneously. Work must be done in order to get these reactions to occur. When endothermic reactions absorb energy, a temperature drop is measured during the reaction. Endothermic reactions are characterized by positive heat flow (into the reaction) and an increase in enthalpy (+ H).

An exothermic reaction is one in which heat is produced as one of the end products. Â Examples of exothermic reactions from our daily life are combustion like the burning of a candle, wood, and neutralization reactions. In an endothermic reaction, the opposite happens. In this reaction, heat is absorbed. Or more exactly, heat is required to complete the reaction. Photosynthesis in plants is a chemical endothermic reaction. In this process, the chloroplasts in the leaves absorb the sunlight. Without sunlight or some other similar source of energy, this reaction cannot be completed.

In exothermic reactions the enthalpy change is always negative while in endothermic reactions the enthalpy change is always positive. This is due to the releasing and absorption of heat energy in the reactions, respectively. The end products are stable in exothermic reactions. The end products of endothermic reactions are less stable. This is due to the weak bonds formed.

‘Endo’ means to absorb and so in endothermic reactions, the energy is absorbed from the external surrounding environment. So the surroundings lose energy and as a result

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the end product has higher energy level than the reactants. Due to this higher energy bonds, the product is less stable. And most of the endothermic reactions are not spontaneous. ‘Exo’ means to give off and so energy is liberated in exothermic reactions. As a result, the surroundings get heated up. And most exothermic reactions are spontaneous.

When we light a matchstick, it is an exothermic reaction. In this reaction, when we strike the stick, stored energy is released as heat spontaneously. And the flame will have lower energy than the heat produced. The energy being released is previously stored in the matchstick and thus do not require any external energy for the reaction to occur.
When ice melts, it will be due to the heat around. The surrounding environment will have a higher temperature than the ice and this heat energy is absorbed by the ice. The stability of the bonds is reduced and as a result and the ice melts into liquid.
Some exothermic reactions in our lives are the digestion of food in our body, combustion reactions, water condensations, bomb explosions, and adding an alkali metal to water.
Reaction rate
Reaction rate, the speed at which a chemical reaction proceeds. It is often expressed in terms of either the concentration (amount per unit volume) of a product that is formed in a unit of time or the concentration of a reactant that is consumed in a unit of time. Alternatively, it may be defined in terms of the amounts of the reactants consumed or products formed in a unit of time. For example, suppose that the balancedchemical equation for a reaction is of the formA + 3B → 2Z.
The rate could be expressed in the following alternative ways:d[Z]/dt, –d[A]/dt,
–d[B]/dt, dz/dt, −da/dt, −db/dtwhere t is the time, [A], [B], and [Z] are the concentrations of the substances, and a, b, and z are their amounts. Note that these six expressions are all different from one another but are simply related. Chemical reactions proceed at vastly different speeds depending on the nature of the reacting substances, the type of chemical transformation, the temperature, and other factors. In general, reactions in which atoms or ions (electrically charged particles) combine occur very rapidly, while those in which covalent bonds(bonds in which atoms share electrons) are broken are much slower. For a given reaction, the speed of the reaction will vary with the temperature, thepressure, and the amounts of reactants present. Reactions usually slow down as time goes on because of the depletion of the reactants. In some cases the addition of a substance that is not itself a reactant, called
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a catalyst, accelerates a reaction. The rate constant, or the specific rate constant, is the proportionality constant in the equation that expresses the relationship between the rate of a chemical reaction and the concentrations of the reacting substances. The measurement and interpretation of reactions constitute the branch of chemistry known as chemical kinetics.

The reaction rate (rate of reaction) or speed of reaction for a reactant or product in a particular reaction is intuitively defined as how fast or slow a reaction takes place. For example, the oxidative rusting of iron under Earth's atmosphere is a slow reaction that can take many years, but the combustion of cellulose in a fire is a reaction that takes place in fractions of a second.
Chemical kinetics is the part of physical chemistry that studies reaction rates. The concepts of chemical kinetics are applied in many disciplines, such as chemical engineering, enzymology and environmental engineering.
Formal definition of reaction rate[edit]
Consider a typical chemical reaction:

a A + b B → p P + q Q

The lowercase letters (a, b, p, and q) represent stoichiometric coefficients, while the capital letters represent the reactants (A and B) and the products (P and Q).

According to IUPAC's Gold Book definition^[1] the reaction rate r for a chemical reaction occurring in a closed system under isochoric conditions, without a build-up of reaction intermediates, is defined as:
where [X] denotes the concentration of the substance X. (Note: The rate of a reaction is always positive. A negative sign is present to indicate the reactant concentration is decreasing.) The IUPAC^[1] recommends that the unit of time should always be the second. In such a case the rate of reaction differs from the rate of increase of concentration of a product P by a constant factor (the reciprocal of its stoichiometric number) and for a reactant A by minus the reciprocal of the stoichiometric number. Reaction rate usually has the units of mol L⁻¹ s⁻¹. It is important to bear in mind that the previous definition is only valid for a single reaction , in a closed system of constant volume. This usually implicit assumption must be stated explicitly, otherwise the definition is incorrect: If water is added to a pot containing salty water, the concentration of salt decreases, although there is no chemical reaction.
For any open system, the full mass balance must be taken into account: in − out + generation − consumption = accumulation

where F_A0 is the inflow rate of A in molecules per second, F_A the outflow, and v is the instantaneous reaction rate of A (in number concentration rather than molar) in a given differential volume, integrated over the entire system volume V at a given moment. When applied to the closed system at constant volume considered previously, this equation reduces to:

where the concentration [A] is related to the number of molecules N_A by [A] = N_A/N₀V. Here N₀ is the Avogadro constant.

For a single reaction in a closed system of varying volume the so -called rate of conversion can be used, in order to avoid handling concentrations. It is defined as the derivative of the extent of reaction with respect to time.

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Here ν_i is the stoichiometric coefficient for substance i, equal to a, b, p , and q in the typical reaction above. Also V is the volume of reaction and C_i is the concentration of substance i.
When side products or reaction intermediates are formed, the IUPAC^[1] recommends the use of the terms rate of appearance and rate of disappearance for products and reactants, properly.
Reaction rates may also be defined on a basis that is not the volume of the reactor. When a catalyst is used the reaction rate may be stated on a catalyst weight (mol g⁻¹ s⁻¹ ) or surface area (mol m⁻² s⁻¹) basis. If the basis is a specific catalyst site that may be rigorously counted by a specified method, the rate is given in units of s⁻¹ and is called a turnover frequency.

Factors influencing rate of reaction[edit]

The nature of the reaction: Some reactions are naturally faster than others. The number of reacting species, their physical state (the particles that form solids move much more slowly than those of gases or those in solution), the complexity of the reaction and other factors can greatly influence the rate of a reaction.
Concentration: Reaction rate increases with concentration, as described by the rate law and explained by collision theory. As reactant concentration increases, the frequencyof collision increases.

Pressure: The rate of gaseous reactions increases with pressure, which is, in fact, equivalent to an increase in concentration of the gas.The reaction rate increases in the direction where there are fewer moles of gas and decreases in the reverse direction. For condensed-phase reactions, the pressure dependence is weak.
Order: The order of the reaction controls how the reactant concentration (or pressure) affects reaction rate.
Temperature: Usually conducting a reaction at a higher temperature delivers more energy into the system and increases the reaction rate by causing more collisions between particles, as explained by collision theory. However, the main reason that temperature increases the rate of reaction is that more of the colliding particles will have the necessary activation energy resulting in more successful collisions (when bonds are formed between reactants). The influence of temperature is described by the Arrhenius equation.

For example, coal burns in a fireplace in the presence of oxygen, but it does not when it is stored at room temperature. The reaction is spontaneous at low and high temperatures but at room temperature its rate is so slow that it is negligible. The increase in temperature, as created by a match, allows the reaction to start and then it heats itself, because it is exothermic. That is valid for many other fuels, such as methane, butane, and hydrogen.

Reaction rates can be independent of temperature (non-Arrhenius) or decrease with increasing temperature (anti-Arrhenius). Reactions without an activation barrier (e.g., someradical reactions), tend to have anti Arrhenius temperature dependence: the rate constant decreases with increasing temperature.

Solvent: Many reactions take place in solution and the properties of the solvent affect the reaction rate. The ionic strength also has an effect on reaction rate.

Electromagnetic radiation and intensity of light: Electromagnetic radiation is

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a form of energy. As such, it may speed up the rate or even make a reaction spontaneous as it provides the particles of the reactants with more energy. This energy is in one way or another stored in the reacting particles (it may break bonds, promote molecules to electronically or vibrationally excited states ...) creating intermediate species that react easily. As the intensity of light increases, the particles absorb more energy and hence the rate of reaction increases.

For example, when methane reacts with chlorine in the dark, the reaction rate is very slow. It can be sped up when the mixture is put under diffused light. In bright sunlight, the reaction is explosive.

A catalyst: The presence of a catalyst increases the reaction rate (in both the forward and reverse reactions) by providing an alternative pathway with a lower activation energy.

For example, platinum catalyzes the combustion of hydrogen with oxygen at room temperature.

Isotopes: The kinetic isotope effect consists in a different reaction rate for the same molecule if it has different isotopes, usually hydrogen isotopes, because of the relative mass difference between hydrogen and deuterium.

Surface Area: In reactions on surfaces, which take place for example during heterogeneous catalysis, the rate of reaction increases as the surface area does. That is because more particles of the solid are exposed and can be hit by reactant molecules.

Stirring: Stirring can have a strong effect on the rate of reaction for heterogeneous reactions.

Diffusion limit: Some reactions are limited by diffusion.

All the factors that affect a reaction rate, except for concentration and reaction order, are taken into account in the reaction rate coefficient (the coefficient in the rate equation of the reaction).
Rate equation[edit]
Main article: Rate equation

For a chemical reaction a A + b B → p P + q Q, the rate equation or rate law is a mathematical expression used in chemical kinetics to link the rate of a reaction to theconcentration of each reactant. It is of the kind:

For gas phase reaction the rate is often alternatively expressed by partial pressures.

In these equations k(T) is the reaction rate coefficient or rate constant, although it is not really a constant, because it includes all the parameters that affect reaction rate, except for concentration, which is explicitly taken into account. Of all the parameters influencing reaction rates, temperature is normally the most important one and is accounted for by the Arrhenius equation.

The exponents n and m are called reaction orders and depend on the reaction mechanism. For elementary (single-step) reactions the order with respect to each reactant is equal to its stoichiometric coefficient. For complex (multistep) reactions, however, this is often not true and the rate equation is determined by the detailed mechanism, as illustrated below for the reaction of H₂ and NO.
For elementary reactions or reaction steps, the order and stoichiometric

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coefficient are both equal to the molecularity or number of molecules participating. For a unimolecular reaction or step the rate is proportional to the concentration of molecules of reactant, so that the rate law is first order. For a bimolecular reaction or step, the number of collisionsis proportional to the product of the two reactant concentrations, or second order. A termolecular step is predicted to be third order, but also very slow as simultaneous collisions of three molecules are rare.

By using the mass balance for the system in which the reaction occurs, an expression for the rate of change in concentration can be derived. For a closed system with constant volume, such an expression can look like

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