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CHEMISTRY REPORT

RATE REACTION

SMA Negeri 2 CIREBONPioneering International School

Jl. Dr. Cipto mangukusumo. 01 Tel. (0231) 203301 - Fax (0231) 239814 Cirebon

TABLE OF CONTENTS :

GOAL ATTEMPT.....................................................................................................................................3

BASIC THEORY........................................................................................................................................3

Reaction Rate..................................................................................................................3

Rate Equation.................................................................................................................3

Molarity..........................................................................................................................5

Dillution of Solution........................................................................................................5

Factors that affects reaction rate....................................................................................6

Collision Theory..............................................................................................................8

TOOLS AND MATERIAL........................................................................................................................9

EXPERIMENTAL PROCEDURES...........................................................................................................9

ORDER REACTION (TABLE)..................................................................................................................9

TABULATION OF DATA......................................................................................................................10

CONCLUSION

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The Goal Attempt :

This experiment has purposed to Determine the order reaction of HCL and Na2S2O3

Basic Theory :

Reaction rate

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 oxidation of iron under the atmosphere is a slow reaction that can take many years, but the combustion of butane in a fire is a reaction that takes place in fractions of a second. The rate of reaction can be defined as molarities per second (M/second).

Rate equation

For a chemical reaction n A + m B → C + D, the rate equation or rate law is a mathematical expression used in chemical kinetics to link the rate of a reaction to the concentration of each reactant. It is of the kind:

In this equation 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 described before, temperature is normally the most important one.The exponents n and m are called reaction orders and depend on the reaction mechanism.

Stoichiometry, molecularity (the actual number of molecules colliding), and reaction order coincide necessarily only in elementary reactions, that is, those reactions that take place in just one step. The reaction equation for elementary reactions coincides with the process taking place at the atomic level, i.e. n molecules of type A colliding with m molecules of type B (n plus m is the molecularity).

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For gases, the rate law can also be expressed in pressure units using, e.g., the ideal gas law.

By combining the rate law with a 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

Examples

For the reaction

The rate equation (or rate expression) is:

The rate equation does not simply reflect the reactants stoichiometric coefficients in the overall reaction: It is first order in H2, although the stoichiometric coefficient is 2 and it is second order in NO.

In chemical kinetics, the overall reaction is usually proposed to occur through a number of elementary steps. Not all of these steps affect the rate of reaction; normally it is only the slowest elementary step that affect the reaction rate. For example, in:

1. (fast equilibrium)2. (slow)3. (fast)

Reactions 1 and 3 are very rapid compared to the second, so it is the slowest reaction that is reflected in the rate equation. The slow step is considered the rate determining step. The orders of the rate equation are those from the rate determining step.

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Molarity

Another way of expressing concentration, the way that we will use most in this course, is called molarity. Molarity is the number of moles of solute dissolved in one liter of solution. The units, therefore are moles per liter, specifically it's moles of solute per liter of solution.

M : Molaritas (mol/lt)

V : Volume (ml)

gram : Massa (gram)

Mr : Massa relative atomic

You must be very careful to distinguish between moles and molarity. "Moles" measures the amount or quantity of material you have; "molarity" measures the concentration of that material. So when you're given a problem or some information that says the concentration of the solution is 0.1 M that means that it has 0.1 mole for every liter of solution; it does not mean that it is 0.1 moles. Please be sure to make that distinction.

Dillution of Solution

Dilution is a reduction in the concentration of a chemical (gas, vapor, solution). It is the process of reducing the concentration of asolute in solution, usually simply by mixing with more solvent. To dilute a solution means to add more solvent without the addition of more solute. The resulting solution is thoroughly mixed so as to ensure that all parts of the solution are identical.

The same direct relationship applies to gases and vapors diluted in air for example. Although, thorough mixing of gases and vapors may not be as easily accomplished.

For example, if there are 10 grams of salt (the solute) dissolved in 1 litre of water (the solvent), this solution has a certain salt concentration/molarity. If

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M =

M = .

one adds 1 litre of water to this solution the salt concentration is reduced. The diluted solution still contains 10 grams of salt/(0.171 moles of NaCl).

Mathematically this relationship can be shown in the equation:

M1 . V1 = M2 . V2

M1 : Molarity of solution before dilution (mol / liter)

V1 : The volume of solution before dilution (ml)

M2 : Molarity of solution after dilution (mol / liter)

V2 : The volume of solution after dilution (ml)

Factors that affects reaction rates

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 influence greatly 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 frequency of 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

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the Arrhenius equation. As a rule of thumb, reaction rates for many reactions double for every 10 degrees Celsius increase in temperature,[2]though the effect of temperature may be very much larger or smaller than this.

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., some radical 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 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, usuallyhydrogen isotopes, because of the mass difference between hydrogen and deuterium.

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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.

All the factors that affect a reaction rate, except for concentration and reaction order, are taken into account in the rate equation of the reaction.

Collision theory

Reaction rate tends to

increase

with concentration phenomenon explained by collision theory

Collision theory is a theory proposed independently by Max Trautz in 1916 and William Lewis in 1918, that qualitatively explains how chemical reactions occur and why reaction ratesdiffer for different reactions. The collision theory can only occur when the suitable particles of the reactant hit with each other. Only a certain percentage of the sum of the collisions cause any noticeable or significant chemical change; these successful changes are called successful collisions. The successful collisions have enough energy, also known as activation energy, at the moment of impact to break the preexisting bonds and form all new bonds. This results in the products of the reaction. Increasing the concentration of the reactant particles or raising the temperature, thus bringing about more collisions and therefore many more successful collisions, increases the rate of reaction.

When a catalyst is involved in the collision between the reactant molecules, less energy is required for the chemical change to take place, and hence more collisions have sufficient energy for reaction to occur. The reaction rate therefore increases. Collision theory is closely related to chemical kinetics.

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Tools and Materials :

Measuring cup : 100 mL Measuring cylinder : 25 mL (2 pieces)

Stopwatch HCl 0,1 M Na2S2O3 0,2 M

Experimental Procedures :

1. Put a cross (x) in the white paper2. Enter the 10 ml of 0.1 M HCl solution to the beaker and place the

beaker on the Cross3. Add 20 ml of 0.1 M Na2S2O3 solution 4. Record the time from the addition to a cross (x) look no further

than the solution5. Repeat the experiment using the numbers 2-4 with Na2S2O3

solution diluted first with water.( as listed in Table)

Order Reaction of Na2S2O3 (Table)

NoVolume

WaktuHCl 0,1 M Penambahan air

terhadap HClTotal Na2S2O3

0,2 MPenambahan air terhadap Na2S2O3

Total

1 10 0 10 20 0 20 108

2 10 0 10 15 5 20 154

3 10 0 10 10 10 20 216

4 10 0 10 5 15 20 268

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Tabulation of Data :

Initial concentration of the reaction Na2S2O3 (M1 . V1 = M2 . V2)

HCl → M1 . V1 = M2 . V2

0,1 . 10 = M2 . 10

M2 = 0,1

Na2S2O3 → i. M1 . V1 = M2 . V2

0,2 . 20 = M2 . 20

M2 = 0,2 M

ii. M1 . V1 = M2 . V2

0,2 . 15 = M2 . 20

3 = M2 . 20

M2 = 320

M2 = 0,15 M

iii. M1 . V1 = M2 . V2

0,2 . 10 = M2 . 20

2 = M2 . 20

M2 = 220

M2 = 0,1 M

iv. M1 . V1 = M2 . V2

0,2 . 5 = M2 . 20

1 = M2 . 20

M2 = 0,05 M

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No HCl Na2S2O3 V

10,1 0,2 1

108

2 0.1 0.15 1154

3 0.1 0,1 1216

4 0.1 0,05 1268

Orde Reaksi Na2S2O3

V 1V 3 = (

0 ,10 ,1

)X (0 ,20 ,1

)y

11081216

= 2y

216108 = 2y

2 = 2y

Y = 1

So, Order reaction of Na2S2O3 is 1

CONCLUSION

Order reaction of HCl is 0

Order reaction of Na2S2O3 is 1

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