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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

displaces H2 gas from water, steam and acids and forms hydroxides

Electrolysis

Potassium



Strontium

Sr 

Calcium

Ca 

Sodium

Na 

Magnesium

Mg 

displaces H2 gas from steam and acids and forms hydroxides

Aluminium

Al 

Carbon



Included for comparison




Manganese

Mn 

displaces H2 gas from steam and acids and forms hydroxides

Smelting with coke

Zinc

Zn 

Chromium

Cr 

Iron

Fe 

displaces H2 gas from acids only and forms hydroxides

Cadmium

Cd 

Cobalt

Co 

Nickel

Ni 

Tin

Sn 

Lead

Pb

Hydrogen gas

H2

included for comparison




Antimony

Sb

combines with O2 to form oxides and cannot displace H2

Heat or physical extraction methods

Arsenic

Ar 

Bismuth

Bi 

Copper

Cu 

Mercury

Hg 

found free in nature, oxides decompose with heating

Silver

Ag 

Paladium

Pd 

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 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 ΔH, or a change in enthalpy. A positive ΔH means that energy is stored and the reaction is endothermic. A negative ΔH 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:

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 (ΔS > 0) of the system. They are denoted by a negative heat flow (heat is lost to the surroundings) and decrease in enthalpy (ΔH < 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 (+ΔH).

An exothermic reaction is one in which heat is produced as one of the end products.   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

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]/dtdz/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 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 engineeringenzymology 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 (abp, 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:

{\displaystyle r=-{\frac {1}{a}}{\frac {d[\mathrm {A} ]}{dt}}=-{\frac {1}{b}}{\frac {d[\mathrm {B} ]}{dt}}={\frac {1}{p}}{\frac {d[\mathrm {P} ]}{dt}}={\frac {1}{q}}{\frac {d[\mathrm {Q} ]}{dt}}}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



{\displaystyle F_{\mathrm {A} 0}-F_{\mathrm {A} }+\int _{0}^{V}v\,dV={\frac {dN_{\mathrm {A} }}{dt}}}where FA0 is the inflow rate of A in molecules per second, FA 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:

{\displaystyle r={\frac {d[A]}{dt}}}where the concentration [A] is related to the number of molecules NA by [A] = NA/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.{\displaystyle r={\frac {d\xi }{dt}}={\frac {1}{\nu _{i}}}{\frac {dn_{i}}{dt}}={\frac {1}{\nu _{i}}}{\frac {d(C_{i}V)}{dt}}={\frac {1}{\nu _{i}}}\left(V{\frac {dC_{i}}{dt}}+C_{i}{\frac {dV}{dt}}\right)}

Here νi is the stoichiometric coefficient for substance i, equal to abp, and q in the typical reaction above. Also V is the volume of reaction and Ci 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.

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