Introduction: Chemistry is a branch of natural science that deals principally with the
properties of substances, the changes they undergo, and the natural laws that
describe these changes. Before moving into the study of chemistry or any other
science, it is important to understand the basis of scientific study, recording and
presenting scientific data. In this regard it makes order for one scientist to
understand the results of a different scientist’s experiments. There must be a
standardized system through which the data can be shared and understood.
There are a number of important concepts associated with recording scientific data and presenting experimental results. Specifically, it is important to know the concepts of units, significant figures, and scientific notation. Units allow scientists to standardize measurements of scientific data, while the rules for expressing significant figures ensure that data is presented honestly and accurately. Scientific notation provides a more convenient method of writing otherwise cumbersome large and small numbers.
When studying chemistry it is important to know and also discuss the more abstract principles of accuracy and precision. Accuracy describes the extent to which an experimental result is correct. Precision refers to the extent to which an experiment can be repeated with the same result. The two concepts are distinctly different, since measurements can be very precise, but not accurate, or accurate but not precise.
INTRODUCTION TO CHEMISTRY
What is chemistry? Chemistry is the study of the interactions of matter with other matter and with energy. However, the definition of chemistry includes a wide range of topics that must be understood to gain a mastery of the topic or even take additional courses in chemistry. Chemistry is sometimes referred to as “the central science” due to its Interconnectedness with a vast array of other STEM disciplines (STEM stands for areas of study in the science, technology, engineering, and math fields).
The definition of chemistry as the study of the interactions of matter with other matter and with energy uses some terms that should also be defined. We start the study of chemistry by defining some of these basic terms.
Chemistry is everywhere and is in everything. Anything we that we touch, smell or see contains one or more chemicals. Many occur naturally but some are man-made.
Chemistry discover natural occurring chemicals and also make new ones never seen before. Chemistry also study the properties of the natural and man-made chemicals. Information about chemistry is used to understand how some chemicals may be modified to make them more useful and they develop the methods to make the modifications.
Chemistry spans the range from qualitative in focus to quantitative. The more
qualitative chemist might work on synthesizing a new compound used in medicine,
the more quantitative work can seem much like physics applied to the microscopic
level of atoms and molecules.
Matter: is defined as anything that occupies space and has mass, and it is all around
us. Solids and liquids are more obviously matter: We can see that they take up space,
and their weight tells us that they have mass. Gases are also matter; if gases did not
take up space, a balloon would stay collapsed rather than inflate when filled with gas.
Solids, liquids, and gases are the three states of matter commonly found on earth
A solid is rigid and possesses a definite shape. A liquid flows and takes the shape
of a container, except that it forms a flat or slightly curved upper surface when acted
upon by gravity. (In zero gravity, liquids assume a spherical shape.) Both liquid and
solid samples have volumes that are very nearly independent of pressure. A gas takes
both the shape and volume of its container.
Figure above shows the three most common states or phases of matter which are
solid, liquid, and gas. A fourth state of matter, plasma, occurs naturally in the
interiors of stars. A plasma is a gaseous state of matter that contains appreciable
numbers of electrically charged particles. The presence of these charged particles
imparts unique properties to plasmas that justify their classification as a state of
matter distinct from gases. In addition to stars, plasmas are found in some other
high-temperature environments (both natural and man-made), such as lightning
strikes, certain television screens, and specialized analytical instruments used to
detect trace amounts of metals. Physical and chemical properties are important
aspects of chemistry. If matter always stayed the same, chemistry would be rather
boring. Fortunately, a major part of chemistry involves change.
A physical change occurs when a sample of matter changes one or more of its
physical properties. For example, a solid may melt or alcohol in a thermometer may
change volume as the temperature changes. However, a physical change does not
affect the chemical composition of matter. A chemical change is the process of
demonstrating a chemical property, such as the burning match. As the matter in the
match burns, its chemical composition changes, and new forms of matter with new
physical properties are created. Note that chemical changes are frequently
accompanied by physical changes, as the new matter will likely have different
physical properties from the original matter.
In Chemistry as a science, a sample of matter that has the same physical and chemical properties throughout is called a substance and in this case, Sometimes the phrase pure substance is used, though the word pure itself isn’t needed. The definition of the term substance is an example of how chemistry has a specific definition for a word that is used in everyday language with a different, vaguer definition. Here, we will use the term substance with its strict chemical definition.
Chemistry recognizes two different types of substances which are elements and compounds. An element can best be defined as the simplest type of chemical substance, it cannot be broken down into simpler chemical substances by ordinary chemical means. There are about 115 elements known to science, of which 80 are stable. The other elements are radioactive. Each element has its own unique set of physical and chemical properties. Examples of elements include iron, carbon, and gold.
A compound is defined as a combination of more than one element. The physical and chemical properties of a compound are different from the physical and chemical properties of its constituent elements that is it behaves as a completely different substance. There are over 50 million compounds known, and more are being discovered daily. Examples of compounds include water, sodium carbonate, and sodium chloride which is the chemical name for common table salt that we use in our daily live to put in our foods.
Elements and compounds are not the only ways in which matter can be present. We frequently encounter objects that are physical combinations of more than one element or compound. Physical combinations of more than one substance are called mixtures.
There are two types of mixtures heterogeneous mixture and homogeneous mixture. A heterogeneous mixture is a mixture composed of two or more substances. It is easy to tell, sometimes by the naked eye, that more than one substance is present whilst homogeneous mixture is a combination of two or more substances that is so intimately mixed that the mixture behaves as a single substance. Another word for a homogeneous mixture is solution. Thus, a combination of salt and steel wool is a heterogeneous mixture because it is easy to see which particles of the matter are salt crystals and which are steel wool. On the other hand, if we take salt crystals and dissolve them in water, it is very difficult to tell that you have more than one substance present just by looking even if you use a powerful microscope. The salt dissolved in water is a homogeneous mixture, or a solution.
Matter can also be described by use of other means especially elements. We can usually divide elements into two as follows, metals and non-metals, and each set shares certain but not always all properties. A metal is an element that is solid at room temperature although mercury is a well-known exception, is shiny and silvery, conducts electricity and heat well, can be pounded into thin sheets a property called malleability, and can be drawn into thin wires a property called ductility. A non-metal is an element that is brittle when solid, does not conduct electricity or heat very well, and cannot be made into thin sheets or wires. Non-metals also exist in a variety of phases and colours at room temperature. Some elements have properties of both metals and non-metals and are called semimetals or metalloids.
The following is a flowchart of the relationships among the different ways of describing matter.
Structure of Atoms: Atoms consist of protons and neutrons in the nucleus, surrounded by electrons that reside in orbitals. Orbitals are classified according to the four quantum numbers that represent any one particular orbital’s energy, shape, orientation, and the spin of the occupying electron. Atomic Structure mainly focus on the electron and the mechanism of describing electrons and their orbitals.
Atomic orbital, an orbital, associated with only one particular atom, in which electrons reside. Though they are called orbitals, atomic orbitals should not be conceived as akin to the orbits of planets rather around a star. Instead, orbitals describe a locus of space in which an electron is likely to reside. Each orbital can hold up to two electrons
Electrons fill up orbitals in a systematic fashion, following the rules of the Aufbau principle. Aufbau principle is German for “building up”, a systematic procedure for determining the electron configuration of any atom. Incorporates the Pauli Exclusion Principle and Hund’s Rule. The configuration of electrons in an atom play a vital role in chemistry. Virtually every chemical process relies on the interactions of electrons between atoms, most particularly on the tendency of atoms to follow the octet rule, the tendency to gain a full valence shell electrons.
Hund’s Rule is a rule which says that, when choosing between orbitals, electrons prefer to go in separate orbitals of the same energy. In this way, every orbital within a particular shell or subshell when the orbitals are not degenerate will be half-filled before any single one orbital becomes completely filled. Degenerate orbitals is orbitals with identical energies. Pauli Exclusion Principle, States that no two electrons in an atom or molecule can have the same set of four quantum numbers.
Quantum Numbers are the four numbers that define each particular electron of an atom. The principle quantum number (n) describes the electrons’ energy and distance from the nucleus. The angular momentum quantum number (l) describes the shape of the orbital in which the electron resides. The magnetic quantum Number (m) describes the orientation of the orbital in space. The spin quantum number describes whether the spin of the electron is positive or negative.
Introduction to chemical Bonding: Chemical reactions involve the making and breaking of bonds. It is essential that we know what bonds are before we can understand any chemical reaction. To understand bonds, we will first describe several of their properties. The bond strength tells us how hard it is to break a bond. Bond lengths give us valuable structural information about the positions of the atomic nuclei. Bond dipoles inform us about the electron distribution around the two bonded atoms. From bond dipoles we may derive electronegativity data useful for predicting the bond dipoles of bonds that may have never been made before.
From these properties of bonds we will see that there are two fundamental types of bonds–covalent and ionic. Covalent bonding represents a situation of about equal sharing of the electrons between nuclei in the bond. Covalent bonds are formed between atoms of approximately equal electronegativity. Because each atom has near equal pull for the electrons in the bond, the electrons are not completely transferred from one atom to another. When the difference in electronegativity between the two atoms in a bond is large, the more electronegative atom can strip an electron off of the less electronegative one to form a negatively charged anion and a positively charged cation. The two ions are held together in an ionic bond because the oppositely charged ions attract each other as described by Coulomb’s Law.
Ionic compounds, when in the solid state, can be described as ionic lattices whose shapes are dictated by the need to place oppositely charged ions close to each other and similarly charged ions as far apart as possible. Though there is some structural diversity in ionic compounds, covalent compounds present us with a world of structural possibilities. From simple linear molecules like H2 to complex chains of atoms like butane (CH3CH2CH2CH3), covalent molecules can take on many shapes. To help decide which shape a polyatomic molecule might prefer we will use Valence Shell Electron Pair Repulsion theory (VSEPR). VSEPR states that electrons like to stay as far away from one another as possible to provide the lowest energy (i.e. most stable) structure for any bonding arrangement. In this way, VSEPR is a powerful tool for predicting the geometries of covalent molecules.
The development of quantum mechanics in the 1920’s and 1930’s has revolutionized our understanding of the chemical bond. It has allowed chemists to advance from the simple picture that covalent and ionic bonding affords to a more complex model based on molecular orbital theory. Molecular orbital theory postulates the existence of a set of molecular orbitals, analogous to atomic orbitals, which are produced by the combination of atomic orbitals on the bonded atoms. From these molecular orbitals we can predict the electron distribution in a bond about the atoms. Molecular orbital theory provides a valuable theoretical complement to the traditional conceptions of ionic and covalent bonding with which we start our analysis of the chemical bond.
What is Stoichiometry: Stoichiometry is at the heart of the production of many things that we use in our daily life. Soap, rubbers, fertilizer, gasoline and deodorant are just a few commodities we use that are chemically engineered, or produced through chemical reactions. All chemically engineered commodities rely on stoichiometry for their production.
The question one has to ask is, what is this so called stoichiometry? The simple response about stoichiometry can be alluded as the calculation of quantities in chemical equations. Given a chemical reaction, stoichiometry tells us what quantity of each reactant we need in order to get enough of our desired product. Because of its real-life applications in chemical engineering as well as research, stoichiometry is one of the most important and fundamental topics in chemistry therefore it cannot be left out as it is the basic understanding point of chemistry and the reactions that occur.
Introduction to the Mole, the mole is the unit of measurement for amount of substance in the international system of units (SI). The unit is defined as the amount of sample of a chemical substance that contains as many constitute particles e.g. atoms, molecules, ions, electrons and protons. The mole is an SI base unit with the unit symbol mol. The mole is widely used in chemistry as a convenient way to express amounts of reactants and products of chemical reactions. For example, the chemical equation 2 H2 + O2 ? 2H2O implies that 2 mol dihydrogen (H2) and 1 mol dioxygen (O2) react to form 2 mol water (H2O). The mole may also be used to represent the number of atoms, ions, or other entities in a given sample of a substance and when it comes to the concentration of a solution, it is commonly expressed by its molarity defined as the amount of dissolved substance per unit volume of solution, for which the unit typically used is moles per litre (mol/l).Giving an example on units of weight makes very interesting, if for instance I have 100 pounds of bowling balls and 100 pounds of feathers, do I have more feathers or more bowling balls? The quantities of feathers and bowling balls would not be equal. An individual feather weighs a lot less than a bowling ball. It would take only about 10 bowling balls to get 100 pounds whereas it would take a lot more feathers.
When you measure quantities in moles, however, the situation is exactly opposite from when you measure according to weight. A mole is defined as the amount of a substance. More specifically, there are 6.02×1023 particles in a mole of substance. Therefore, if you had 1 mole of feathers and 1 mole of bowling balls, you would have 6.02×1023 feathers and 6.02×1023 bowling balls. Now suppose you were asked the question, “Which weighs more, 100 moles of feathers or 100 moles of bowling balls?” The answer this time would be the bowling balls. Although there is an equal number of both feathers and bowling balls, an individual bowling ball weighs much more than an individual feather, and so an equal number of bowling balls would weigh more than an equal number of feathers.
Now, let’s return to the number 6.02×1023. This number is known as Avogadro’s number and you should definitely commit it to memory. You are probably wondering why it’s so large, and it does indeed look intimidating without the exponential notation: 6.02 ×1023 = 60,200,000,000,000,000,000,000,000
Although we can never have a mole of bowling balls, we can still be calculating
moles of compounds, molecules, atoms, and ions. These representative particles are
extremely and incredibly small. That is why there are so many particles in a mole of
substance. When we appreciate just how small these particles are, 6.02×1023 stops
Seeming like such a crazy number.
Gases: The first step to understanding gases is to spell out what exactly a gas is. Gases have two properties that set them apart from solids and liquids. First, gases spontaneously expand to fill the container they occupy, no matter its size. In other words, a gas has no fixed volume or shape. Secondly, gases are easily compressible.
In an imaginable situation such that a gas as a busy swarm of molecules. Each molecule moves randomly and travels great distances before bouncing off another molecule. This occurs because the individual molecules comprising a gas are generally far apart. In fact, for a gas at low pressure, we can approximate that aside from a few random collisions, individual gas molecules do not interact. This approximation is what separates gases from solids and liquids, whose molecules always interact. The series on Gases should therefore seek to use this approximation about gases to establish the ideal gas law and the kinetic molecular theory. The ideal gas law macroscopically describes how gases behave under nearly all conditions and the kinetic molecular theory describes how sub-microscopic gas molecules interact with each other.
Pressure: Of the three general terms used to describe gases (volume, temperature, pressure), pressure is the least familiar. Before we can delve into the gas theories, we need a firm understanding of it. In general, Pressure is defined as the force exerted on a given area: P =
Note that pressure is directly proportional to force and inversely proportional to area.
Thus, pressure can be increased either by increasing the amount of force or by
decreasing the area over which it is applied; pressure can be decreased by
decreasing the force or increasing the area. Ice skates are familiar examples of the
effects of pressure. The area of the blades of a skate are much smaller than, say,
the soles of our feet. So if we strap on ice skates, our weight will act on an area
much smaller than it would if we were wearing normal shoes. Since A decreases
while F stays the same, by @@[email protected]@, the pressure we exert on the ice will
be much greater if we are wearing skates. This pressure is often enough to melt a
layer of ice, which allows our skate to glide smoothly across an ice rink. If we try the
same manoeuvre with normal shoes, we will not generate enough pressure to melt
the ice and won’t get anywhere fast. So how does pressure relate to gases? If we
can remember that, a gas will fill any container that holds it. It is easy to see why with
our swarm analogy. If a compact swarm of molecules is placed into a large
container, the individual molecules will move about randomly and eventually stray
from their original dimensions. Eventually, some intrepid molecules will reach the
walls of the container. When they do, they will impact the walls of the container.
These impacts generate a force, and hence a pressure on the walls of the container.
Introduction to chemical solution: Solutions are homogeneous mixtures of two or more pure substances. The component that is in the largest amount is the solvent and the minor component is the solute. In a solution, the solute is dispersed uniformly throughout the solvent. Addition of solutes to a solvent changes the properties of the liquid, if we are to give an example that antifreeze is added to car radiators to prevent the coolant from freezing during very cold seasons and boiling in the hot seasons. It is therefore important to understand the properties of solutions before we can even begin to understand those reactions. Perhaps the most salient characteristic of a solution is its concentration, how much solute is dissolved in what amount of solvent. Several different units of concentration like mass percent, mole fraction, molarity, normality, and molality have been developed for use in different situations. Solution composition. explains the definitions and uses of those units and why it is necessary to have so many different units of concentration.
After we have discussed the units of concentration, we explore the questions of why solutions form at all and what factors affect the solubility of solutes in different solvents. As we learn and see, like dissolves like. Non-polar solvents dissolve non-polar solutes better than polar solvents and polar solvents dissolve polar solutes better than non-polar solvents. Raising the temperature of a solution will increase the solubility of most solid solutes. Likewise, according to Henry’s law, increasing the pressure above a solution will increase the solubility of most gaseous solutes. There are several others Properties of a solution that depend only on the concentration of solute particles and are called colligative properties. They include changes in the vapour pressure, boiling point, and freezing point of the solvent in the solution. The magnitudes of these properties depend only on the total concentration of solute particles in solution, not on the type of particles. The total concentration of solute particles in a solution also determines its osmotic pressure. This is the pressure that must be applied to the solution to prevent diffusion of molecules of pure solvent through a semipermeable membrane into the solution. Ionic compounds may not completely dissociate in solution due to activity effects, in which case observed colligative effects may be less than predicted. Each of the properties is discussed in details in colligative properties.
Colligative Properties: The properties of solutions differ from those of pure substances that are used to form them. The physical properties that are affected are vapour, pressure, boiling and freezing and osmotic pressure. For example, water has a freezing point of 00C and the boiling point of 1000C. When appropriate solute is introduced into water making an aqueous, the freezing point is lowered while the boiling point is elevated. The changes in the freezing and boiling point depend only on the number of particles present known as colligative properties. It is therefore important to note that all solutions with the same concentration example molarity contain the same number of particles. Let say dissolving 1M of glucose will give 1M of concentration of glucose particles since it no ionic but 1M of NaCl or Cacl2 will not give a concentration of 1M since NaCl dissolve to give two ions while cacl2 dissociates to give three ions respectively. The sum of the concentration of the dissolved ions dictates the physical properties.
Solution Process: In order for a salute to be dissolved in a solvent, the attractive
forces between the solute and solvent particles must be great enough to overcome the
attractive forces within the pure solvent and pure solute. The solute and the solvent
molecules in a solution are expanded compared to their position within the pure
The process of expansion, for both the solute and solvent, involves a change in the
energy of the system: this process can be either exothermic or endothermic. After
dissolving, the solute is said to be fully solvated usually by dipole-dipole or ion-
dipole forces, and when the solvent is water, the solute is said to be hydrated. The
separation of the solute particles from one another prior to dissolving is an
endothermic process for both solvent and solute, solute (steps 1 and 2), but when the
solute and solvent combine with each other, this is an exothermic process (step 3). If
the energy released in step 3 is greater than the energy absorbed in steps 1 and 2, the
solution forms and is stable. The term solubility refers to the maximum amount of
material that will dissolve in a given amount of solvent at a given temperature to
produce a stable solution.
Introduction to acids and bases: Acids and bases play a central role in chemistry because, with the exception of redox reactions, every chemical reaction can be classified as an acid-base reaction. Our understanding of chemical reactions as acid-base interactions comes from the wide acceptance of the Lewis definition of acids and bases, which supplanted both the earlier Bronsted-Lowry concept and the first definition, the Arrhenius model. Arrhenius first defined acids as proton (H+) producers in aqueous solution and bases as hydroxide (OH-) producers. Although this model is intuitively correct, it is limited to substances that include proton and hydroxide groups. Bronsted and Lowry proposed the more general definitions of acids and bases as proton donors and acceptors, respectively. Unlike the Arrhenius conception, the Bronsted-Lowry model accounts for acids in solvents other than water, where the proton transfers do not necessarily involve hydroxide ions. But the Bronsted-Lowry model fails to explain the observation that metal ions make water more acidic (discussed in calculating pH’s). Finally, Lewis gave us the more general definition of acids and bases that we use today. According to Lewis, acids are electron pair acceptors and bases are electron pair donors. Any chemical reaction that can be represented as a simple exchange of valence electron pairs to break and form bonds is therefore an acid-base reaction.
Acid-base chemistry is important to us on a practical level as well, outside of laboratory chemical reactions. Our bodily functions, ranging from the microscopic transport of ions across nerve cell membranes to the macroscopic acidic digestion of food in the stomach, are all ruled by the principles of acid-base chemistry. Homeostasis, the temperature and chemical balances in our bodies, is maintained by acid-base reactions. For example, fluctuations in the pH, or concentration of hydrogen ions, of our blood is moderated at a comfortable level through use of buffers. Learning how buffers work and what their limitations are can help us to better understand our physiology. We will start by introducing fundamentals of acid-base chemistry and the calculation of pH, and then we will cover techniques for measuring pH. We learn about buffers and see how they are applied to measure the acidic content of solutions through titration.
Introduction to electrochemistry: Electrochemistry is the study of the exchange between electrical and chemical energy, it has important applications in everyday life stretching from the battery that powers our portable radio to the electro refining that produces the copper pipes carrying our drinking water. Those electrochemical processes utilize oxidation and reduction reactions. An oxidation involves the loss of one or more electrons from a chemical species while a reduction is the gain of one or more electrons by a chemical species. When an oxidation and a reduction are paired together in a redox reaction, electrons can flow from the oxidized species, the reducing agent or reductant, to the reduced species, the oxidizing agent or oxidant. That electron flow can either be spontaneously produced by the reaction and converted into electricity, as in a galvanic cell, or it can be imposed by an outside source to make a non-spontaneous reaction proceed, as in an electrolytic cell.
Introduction to reaction kinetics: Kinetics, the study of the rates of chemical reactions, has a profound impact on our daily lives. Even though some reactions are thermodynamically favourable, such as the conversion of diamonds into graphite, they do not occur at a measurable rate at room temperature. Other reactions, like the explosive reaction between vinegar and baking soda, occur almost instantaneously. Imagine a world where all thermodynamically favoured processes occurred at the same rate–life could not exist under such circumstances because biological processes rely on the kinetic stability of many unstable compounds. Kinetics answers questions about rate, how fast reactions go, and mechanisms, the paths molecules take in going from reactants to products.
To describe the rate of a reaction, it is important to derive the rate law for a chemical reaction and discuss the factors affecting rate. Additionally, describing the experimental techniques, such as the method of initial rates and fitting data to plots based on the integrated rate law is used to determine the rate law for an unknown reaction.
It is also important to have a mechanism and discuss how to determine the path a reaction takes by analyzing and predicting the series of elementary steps that comprise it. By comparing the rate law for a proposed mechanism and other mechanistic predictions to experimental data, we can test the validity of a mechanism. Mechanisms can never be proven exactly, but we can rule out mechanisms that disagree with experimental observations. It is therefore important to use reaction coordinate diagrams to understand and to visualize reaction mechanisms, thermodynamics, and activation energies. Catalysts and intermediates can be important factors in reaction mechanisms, and they provide interesting examples of mechanism problems.
Thermodynamics: is the study of thermal, electrical, chemical and mechanical forms of energy. This study of thermodynamics crosses many disciplines including physics, engineering and chemistry. Of the various branches of thermodynamics, the most important to chemistry is the study of change in energy during a chemical reaction. Thermodynamics plays an important role in our understanding of electrochemical processes. It can tell us whether a given redox reaction is spontaneous and therefore whether it is able to provide useful electrical energy. Thermodynamics also describes how to add reduction potentials to determine the cell potential for a galvanic or electrolytic cell. It also permits us to add reduction potentials of two reductions to calculate the potential of a new, or even theoretical, reduction half-reaction. Thermodynamics provides further insight into electrochemical cells at non-standard conditions in its derivation of the Nernst Equation. The Nernst Equation allows us to calculate the cell potential at any conditions and suggests the construction of concentration cells such as pH meters or other ion-selective electrodes.
Combining ideas from thermodynamics, stoichiometry, and basic electrical theory makes this section the most important to understand if you wish to become proficient at doing electrochemistry problems. Such problems tend to be among the more difficult ones in an exam because they cut across many fields in chemistry and require deep analytical thought and a thorough understanding of electrochemical cells. Here, we get to know and love formulae as ?Go = -nFEo which is useful for converting free energy and potentials and E = Eo – (RT/nF) ln Q the Nernst Equation.
Despite all that thermodynamics has to offer to electrochemistry, it requires much algebra and clear logic. Most of all, it is about understand the big picture as well as how to do the present problems. Chemistry seeks to study the natural world but also seek to improve it by modification on a molecular level. Because everything is a chemical, and chemistry is said to be one of the foundations of modern industrial economies. It Improves our World
Advancements in the field of chemistry have brought about major improvements in
our world. Improvements range from new medicines that cure disease, to new
materials that make us safer and stronger, to new sources of energy that enable new
In concluding my essay, I therefore state that I have learnt that, Chemistry is a science of studying about the interactions of matter with other matter and with energy uses. Matter has been defined as anything that has mass and takes up space giving computer, food and dirt as some of the examples. Structure of atoms is another part that has been studied and learnt that Atoms consist of protons and neutrons in the nucleus, surrounded by electrons that reside in orbitals. Chemical reactions involving the making and breaking of bonds has also been learned as is essential to know what bonds are in order to understand any chemical reaction by first describing several of their properties.
Stoichiometry is another important aspect of learning chemistry which should not be left out in order to have a wider understanding of basic chemistry and can be described as the heart of chemical reactions and production because all chemically engineered commodities rely on stoichiometry for their production. Properties of solutions and describing a solution as homogenous mixture of two or more substance exit in a single phase and referring a mixture where water is a solvent as an aqueous solution.
Kinetics as the study of the rates of chemical reactions, though some reactions are thermodynamically such as the conversion of diamonds into graphite, other reactions, like the explosive reaction between vinegar and baking soda. The mole as the unit of measurement for amount of substances with the unit symbol mol and also learnt about gasses, pressure acid and bases.
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