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Introduction to Chemistry Study Guide

Chemical Level of Organization

Introduction

Chemistry

Chemistry is the study of the composition, structure, and properties of substances, as well as their reactions with one another.

Matter

All living and nonliving things are made of matter .

Matter – anything that takes up space and has mass.

Elements

The matter of the universe is composed of a limited number of basic substances called elements.

Element – a substance that cannot be decomposed into simpler substances by ordinary chemical reactions.

An element is a substance that consists only of atoms with the same number of protons (designated the atomic number), and therefore the same nuclear charge.

Common examples of elements include carbon, hydrogen, iron, sodium, and chlorine.

There are a total of 118 chemical elements. The first 94 of these elements are believed to occur naturally on the earth. The elements with atomic numbers 95 and above do not occur naturally, and are known only as a result of their synthesis in nuclear reactors or particle accelerators.

Chemical Symbols

Each element is designated by a chemical symbol of either one or two letters that stands for its English or Latin name.

Thus H is the symbol for hydrogen, O for Oxygen, C for carbon, Cl for chlorine, Mg for magnesium, K for potassium (the Latin name is Kalium), Na for sodium (the Latin name is Natrium), etc.

Exercise

For each of the symbols, give the name of the element

Symbol Element

N ____________________

P ____________________

S ____________________

Ca ____________________

Mg ____________________

For each element, give the chemical symbol

Element Symbol

Iron ____________________

Iodine ____________________

Mercury ____________________

Lead ____________________

Flourine ____________________

Elements important in Living Organisms

The most abundant elements in living organisms are carbon, hydrogen, oxygen, nitrogen, phosphorus, and sulfur.

Other elements such as calcium, sodium, potassium, magnesium, and iron are not as abundant but are still essential for life.

Trace elements are present in minute quantities yet are also essential for life. Examples include manganese, zinc, copper, and iodine.

Atoms

Matter is composed of tiny units known as atoms.

Atom – the smallest unit of an element, not divisible by ordinary chemical means.

A chemical symbol represents one atom of the element, e.g. N stands for a single atom of nitrogen.

Atomic Structure

Atoms are composed of three fundamental particles: protons, neutrons, and electrons.

Fundamental Particles

Particle Mass Charge

(µ or Daltons) (electronic charge units)

Electron 5.485 7990 x 10-4 ‒1

Proton 1.007 276 47 +1

Neutron 1.008 664 90 0

The protons and neutrons are found in the nucleus of the atom. The nucleus contains the positive charge and almost all the mass of an atom. Each proton carries an electronic charge of +1. The neutrons, as their name implies, have no charge. Protons and neutrons have roughly the same mass, which is close to one.

Atomic Number

Atomic number – the number of protons in the nucleus of an atom.

The number of protons in the nucleus is unique for each element.

The atomic number is usually written as a subscript immediately before the chemical symbol.

For example, 1H indicates that the atomic number of hydrogen is one, i.e. its nucleus contains only one proton. Similarly, 😯 indicates that oxygen nuclei contain eight protons.

Mass Number

The mass number is the total number of protons and neutrons in a nucleus.

The mass number is commonly written as a superscript before the chemical symbol.

For example, most atoms of oxygen contain eight protons and eight neutrons; the mass number is therefore 16 and the nucleus can be symbolized as O16 or, if we wish to show both the atomic number and the mass number as 8O16.

Other Examples:

17Cl35 Chlorine 17 protons 18 neutrons

11Na23 Sodium 11 protons 12 neutrons

6C12 Carbon 6 protons 6 neutrons

7N14 Nitrogen 7 protons 7 neutrons

15P31 Phosphorous 15 protons 16 neutrons

If we are given the atomic number and the mass number, we can determine the number of protons, the number of neutrons, and the number of electrons in a normal, neutral atom. For example in 17Cl35 the atomic number is 17 so there are 17 protons. The mass number, the total number of protons plus neutrons is 35. If we subtract the atomic number (17) from the mass number (35) we get the number of neutrons (18). We can also determine the number of electrons. In a normal neutral atom, the number of electrons equals the number of protons. So in chlorine, if there are 17 protons there are also 17 electrons.

Isotopes

Isotopes – forms of the same element which have the same number of protons, but which differ in the number of neutrons in the nucleus.

All atoms of a particular element must have the same number of protons. If the number of protons changed, the atom would become a different element. However, the number of neutrons can be different for atoms of the same element. These different forms of the same element are known as isotopes. Because the number of neutrons is different for the isotopes of an element, so also is the mass number.

Examples:

There are three different forms of the element Hydrogen:

Hydrogen 1 Hydrogen 2 Hydrogen 3

1H1 1H2 1H3

Hydrogen 1 has one proton in the nucleus and no neutrons, Hydrogen 2 has one proton and 1 neutron, and Hydrogen 3 has 1 proton and 2 neutrons.

There are three isotopes of Oxygen:

Oxygen 16 Oxygen 17 Oxygen 18

8O16 8O17 8O18

Oxygen 16 has 8 protons and 8 neutrons, Oxygen 17 has 8 protons and 9 neutrons, and Oxygen 18 has 8 protons and 10 neutrons.

There are three naturally occurring isotopes of carbon: carbon 12, carbon 13, and carbon 14.

Carbon 12 Carbon 13 Carbon 14

6C12 6C13 6C14

It is possible to create new isotopes that do not exist in nature using nuclear reactors.

Some isotopes are radioactive. Their nuclei are unstable and tend to break down, emitting energy in the form of radiation.

Importance of Radioisotopes

Radioisotopes have been used as tracers to identify the steps in metabolic reactions. For example, the reactions of photosynthesis were worked out using radioactive carbon dioxide.

Radiation is used to treat cancer.

Radioisotopes are used to visualize structures in the body to locate disorders. For example, the coronary arteries of the heart can be made visible, allowing the detection of blockage.

The Electrons

The electrons are negatively charged particles that encircle the nucleus of the atom. Electrons have very little mass.

Each electron has a charge of –1; its charge is exactly opposite to that of a proton.

In a normal neutral atom, the number of electrons is equal to the number of protons.

The positive charges of the protons cancel out the negative charges of the electrons, making the atom as a whole electrically neutral.

Consequently, in a neutral atom, the atomic number represents both the number of protons inside the nucleus and the number of electrons circling around the nucleus.

Examples:

17Cl35 17 protons 18 neutrons 17 electrons

19K39 19 protons 20 neutrons 19 electrons

Orbitals

The electrons travel around the nucleus of the atom in regions known as orbitals.

The distance of an electron from the nucleus is a function of its energy; the higher the energy, the farther from the nucleus will be the probable location of the electron.

The average energy levels of electrons in an atom correspond to a series of so-called electron shells, which can conveniently be represented by concentric circles located at specified distances from the nucleus.

In an atom of oxygen, for example, there are two electrons in the first shell and six in the second shell.

The Orbitals of electrons may have different shapes. In the first electron shell, this shape is always spherical (it is symbolized by s). In the second electron shell, both the spherical shape (s) and a dumbbell shape (symbolized by p) occur. Additional shapes occur in succeeding electron shells.

It has been shown that there is a maximum number of electrons that each shell can contain. The first electron shell can contain a maximum of 2 electrons, the second shell can contain 8, the third shell 18, and the fourth shell 32, etc.

Although the third and successive shells can hold more than eight electrons, they are in a particularly stable configuration when they contain only eight. For our purposes, then, the first shell can be considered complete when it holds two electrons and every other shell can be considered complete when it holds eight electrons.

Electron Distribution and the Chemical Properties of Elements

If the outer shell of electrons is complete, as it is, for example in Helium, which contains 2 electrons in its outer shell, or in neon, which contains 8 electrons in its outer shell, the element has very little tendency to react chemically with other atoms.

Valence electrons – the electrons in the outer shell of an atom are known as valence electrons.

The valance electrons are important in determining the chemical properties of elements and whether they will combine with one another.

Electron Shell Diagrams

We can represent the structure of atoms using Electron Shell or Bohr diagrams. In such a diagram, the number of protons and neutrons are indicated in a circle that represents the nucleus of the atom. The electrons are placed in Orbitals around the nucleus.

Exercise

Given the information below, construct Bohr diagrams of each atom.

19K39 12Mg24 15P31

Atomic Mass

The atomic mass of an element is the ratio of its mass to one twelfth the mass of an atom of carbon-12, a unit known as a Dalton or µ. The atomic mass is calculated by averaging the atomic masses of all the chemical element’s isotopes, weighed by isotopic abundance and dividing it by one Dalton (µ), which is equal to 1.660538782 x 10−27 kg.

Chemical Bonds

Elements combine to form molecules and compounds.

Molecule – consists of two or more atoms that have been bound together by chemical bonds

A molecule is the smallest chemical unit of a substance that is capable of a stable, independent, existence.

Compound – a compound is composed of two or more different kinds of elements joined together by chemical bonds.

The difference between a molecule and a compound is that in a molecule the elements comprising the molecule may be the same or different. For example, H2, N2, O2, CH4, and C6H12O6 are all molecules. In the above list, CH4, and C6H12O6 also are compounds because they contain different kinds of elements. Water (H2O) is also a compound. However, H2, N2, and O2 are molecules of the element hydrogen, nitrogen, and oxygen respectively. They contain only one kind of element.

Elements combine to form compounds.

Example:

Sodium + Chlorine → Sodium Chloride

There are two types of Chemical Bonds: Ionic bonds and Covalent bonds.

Ionic bonding – bonding by transfer of electrons.

Covalent bonding – bonding by sharing of electrons.

When atoms of elements combine, the atoms usually become more stable by completing their outer shells of electrons (2 electrons for the first shell, eight electrons for the outer shells).

Ionic Bonding

Ionic bonding is bonding by transfer of electrons. Electrons are transferred from the outer shell of one atom to the outer shell of a second atom. By this process both atoms usually attain stability by completely filling their outer shells with electrons.

An example of ionic bonding is the reaction in which sodium combines with chlorine to form sodium chloride.

Sodium + Chlorine → Sodium Chloride

The substances on the left side of the equation, sodium and chlorine, are reactants. The arrow means yields. The substance(s) on the right side of the equation are products. The equation reads: sodium and chlorine react to yield sodium chloride.

Using chemical symbols,

11Na23 + 17Cl35 → NaCl

Using Bohr diagrams,

Examining the Bohr diagrams for sodium and chlorine, we can see that the sodium atom contains 11 electrons. Two electrons are in the first shell, eight in the second, and there is 1 electron in the outermost shell. Sodium would attain greater stability if it had a complete outer shell of 8 electrons. Chlorine has 17 electrons. Two electrons are in the first shell, 8 electrons are in the second shell, and the remaining 7 electrons are in the third and outermost shell. Chlorine would attain a more stable arrangement by having a complete outer shell of eight electrons. Both sodium and chlorine could attain complete outer shells containing eight electrons if the single electron in the outer shell of sodium were transferred to the outer shell of chlorine.

When sodium combines with chlorine, an electron is transferred from the outer shell of sodium to the outer shell of chlorine. This creates an ionic bond joining the two elements together.

When sodium loses an electron by transferring it to chlorine, the sodium atom becomes electrically charged. Prior to the transfer, the sodium ion was electrically neutral. There were 11 protons and therefore 11 positive charges in the nucleus. 11 electrons with 11 negative charges canceled these. When the electron was transferred from sodium to chlorine, the electrical charges no longer equaled each other. There are 11 positive charges contributed by the 11 protons, but now there are only 10 negative charges from the 10 electrons. The sodium now has become a charged particle – a sodium ion .

Ion – a charged particle. A charged atom or group of atoms.

By reacting with sodium, the chlorine gains an electron. It now has 18 electrons and 17 protons. This gives it a charge of –1. The chlorine atom is now a chloride ion.

A useful principle to learn is that opposite electrical charges attract one another. Like charges repel one another.

Opposite electrical charges attract one another. The positively charged sodium ion is attracted to the negatively charged chlorine ion. This creates an ionic bond joining the two ions together.

The product of the reaction is sodium chloride, that is, ordinary table salt.

Positive ions are called cations .

Negative ions are called anions .

These ions are named according to their migration in an electrical field. If the ions are placed in an electrolytic cell and are in solution, the positive ions will migrate toward the negative terminal or cathode. They are therefore called cations. The negative ions will migrate toward the positive terminal or anode. They are therefore called anions.

Covalent Bonding

Covalent bonding is bonding by sharing of electrons. In forming a single covalent bond, two atoms each share one of their electrons with the other. These two shared electrons effectively fill an orbital in each atom and thus form a covalent bond between these two atoms.

The atoms of the elementary gases hydrogen, oxygen, nitrogen, fluorine, and chlorine, form stable diatomic (consisting of two atoms) molecules by covalent bonding.

A molecule is the smallest chemical unit of a substance that is capable of a stable, independent existence.

Formation of H2

Two hydrogen atoms join together to form a molecule of hydrogen gas. When the atoms combine, each hydrogen atom shares one of its electrons with the other atom. Each hydrogen atom has one electron circling its nucleus. Its outer shell of electrons would be complete if it had two. In order to achieve this stable state, each hydrogen atom shares its electron with the other. Most of the time, the single electron will orbit around its own nucleus. However, part of the time, it will circle the nucleus of the other atom. This creates a strong covalent bond that holds the two atoms together. The sharing of the electrons produces the effect of complete outer shells containing two electrons. For at least some of the time, the two electrons may orbit one of the nuclei, giving it in effect, two electrons. At other times they will circle the other nucleus. In this way they attain stable outer shells.

Formation of water

The oxygen atom has six electrons in its outer shell. It would attain greater stability by gaining two electrons, thereby completing its outer shell with eight electrons. In forming a molecule of water, the oxygen atom combines with two atoms of hydrogen. Each of the hydrogen atoms shares one of its electrons with oxygen; together with the six electrons that oxygen already has, this gives oxygen a complete outer shell at least part of the time. Each hydrogen atom borrows one electron from oxygen, giving each a complete outer shell as these electrons orbit the hydrogen nuclei. By sharing two pairs of electrons the two hydrogen atoms and the oxygen atom are joined by covalent bonds. The electrons are not shared equally between the hydrogen atoms and oxygen atom. Because oxygen is a larger atom, it has a greater electronegativity. This means that it has a greater attraction for the electrons than hydrogen. As a result, the electrons spend more time circling around the oxygen than the hydrogen. This leads to a greater negative charge around the oxygen, and because the negatively charged electrons are pulled away from the positively charged hydrogen nuclei, a positive charge develops around the hydrogens. Water, thus, is a polar molecule i.e. a molecule with an unequal distribution of electrical charge.

Weak chemical bonds

Hydrogen Bonds

Hydrogen bond – a weak chemical bond between a negatively charged nitrogen or oxygen atom and a positively charged hydrogen atom.

Because opposite electrical charges attract one another, a negatively charged oxygen atom will be attracted to a positively charged hydrogen atom. This creates a weak chemical bond between the two atoms. For example, if two water molecules are adjacent to one another, the negative oxygen end of one water molecule is attracted to the positive hydrogen end of the other water molecule. These bonds create links between water molecules and are responsible for the forces of cohesion observed between water molecules. Another example of hydrogen bonds is found in DNA. In DNA, nitrogenous bases are stacked in pairs in the center of the DNA molecule between the two strands which comprise the molecule. Hydrogen bonds form between the negatively charged nitrogen atoms of nitrogenous bases and positively hydrogen atoms of adjacent nitrogenous bases. This creates hydrogen bonds that hold the two strands of DNA together.

Chemical reactions

Synthesis Reactions

When two or more atoms, ions, or molecules combine to form larger molecules, the process is called a synthesis reaction.

Example:

2H2 + O2 → 2H2O

Decomposition Reactions

In a decomposition reaction, large molecules are broken down to form smaller ones.

Example:

C6H12O6 + O2 → CO2 + H2O

Exchange Reactions

In an exchange reaction, the atoms of one compound switch places with atoms in another compound.

HCl + NaHCO3 ⇌ NaCl + H2CO3

Reversible reactions

The above reaction is also a reversible reaction. It can go in either direction. In the above reaction HCl can combine with NaHCO3 to form NaCl and H2CO3, or NaCl can combine with H2CO3 to form HCl and NaHCO3.

Chemical Compounds

Compounds can be divided into inorganic compounds and organic compounds.

Inorganic Compounds

Inorganic compounds – compounds which do not contain carbon.

Carbon dioxide is an exception to the above definition. Carbon dioxide contains carbon. Nevertheless, it is considered to be an inorganic compound. This is because the properties of carbon dioxide are more like those of the other inorganic compounds.

Inorganic compounds are usually small, ionically bonded molecules. They include water, many salts, such as NaCl, and many acids, such as HCl.

Organic Compounds

Organic compounds – compounds which contain carbon.

Carbon dioxide is an exception to the above definition also. Although it contains carbon, It is classified as an inorganic rather than organic compound.

Organic compounds are held together primarily by covalent bonds. They tend to be very large molecules. Organic compounds include carbohydrates, lipids, proteins, and nuclei acids.

References

McQuarrie D, Rock PA, Gallogly EB. General Chemistry. 4th ed. Mill Valley (CA): University Science Books; 2011.

Introduction to Chemistry Study Guide draft 9 (Corrected 12/26/2018)

PAGE

1

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Chapter 2 Lecture Outline

Understanding Biology

THIRD EDITION

Kenneth A. Mason

Tod Duncan

Jonathan B. Losos

© 2021 McGraw Hill. All rights reserved. Authorized only for instructor use in the classroom.

No reproduction or further distribution permitted without the prior written consent of McGraw Hill.

Because learning changes everything.®

The Nature of Molecules and the Properties of Water

Chapter 2

©Fuse/Gett y Images

© McGraw Hill

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All Matter is Composed of Atoms

Matter has mass and occupies space

All matter is composed of atoms

Atoms are composed of subatomic particles

© McGraw Hill

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3

Atomic Structure

Atoms are composed of three types of subatomic particles

Protons

Positively charged particles

Located in the nucleus

Neutrons

Neutral particles

Located in the nucleus

Electrons

Negatively charged particles

Found in orbitals surrounding the nucleus

© McGraw Hill

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4

Figure 2.2

© McGraw Hill

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Atomic number

Number of protons equals number of electrons

Atoms are electrically neutral

Atomic number = number of protons

Every atom of a particular element has the same number of protons

Element

Any substance that cannot be broken down to any other substance by ordinary chemical means

© McGraw Hill

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Atomic mass

Mass or weight?

Mass – refers to amount of substance

Weight – refers to the force gravity exerts on a substance

Sum of protons and neutrons is the atom’s atomic mass

Each proton and neutron has a mass of approximately 1 Dalton

© McGraw Hill

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Electrons

Negatively charged particles located in orbitals

Neutral atoms have same number of electrons and protons

Ions are charged particles – unbalanced

Cation – more protons than electrons = net positive charge

Anion – fewer protons than electrons = net negative charge

© McGraw Hill

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Isotopes

Atoms of a single element that possess different numbers of neutrons

Radioactive isotopes are unstable and emit radiation as the nucleus breaks up

Half-life – time it takes for one-half of the atoms in a sample to decay

© McGraw Hill

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Figure 2.3

© McGraw Hill

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Electron arrangement

Key to the chemical behavior of an atom lies in the number and arrangement of its electrons in their orbitals

Bohr model – electrons in discrete orbits

Modern physics defines orbital as area around a nucleus where an electron is most likely to be found

No orbital can contain more than two electrons

© McGraw Hill

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Figure 2.4

© McGraw Hill

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Atomic energy levels

Electrons have potential energy related to their position

Electrons farther from nucleus have more energy

Be careful not to confuse energy levels, which are drawn as rings to indicate an electron’s energy, with orbitals, which have a variety of three-dimensional shapes and indicate an electron’s most likely location

© McGraw Hill

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Figure 2.5

© McGraw Hill

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Redox

During some chemical reactions, electrons can be transferred from one atom to another

Still retain the energy of their position in the atom

Oxidation = loss of an electron

Reduction = gain of an electron

© McGraw Hill

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Elements

Periodic table displays elements according to valence electrons

Valence electrons – number of electrons in outermost energy level

Inert (nonreactive) elements have all eight electrons

Octet rule – atoms tend to establish completely full outer energy levels

© McGraw Hill

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Periodic Table of the Elements

© McGraw Hill

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Figure 2.6b

90 naturally occurring elements

Only 12 elements are found in living organisms in substantial amounts

Four elements make up 96.3% of human body weight

Carbon, hydrogen, oxygen,

Organic molecules contain primarily CHON

Some trace elements are very important

© McGraw Hill

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Chemical Bonds

Molecules are groups of atoms held together in a stable association

Compounds are molecules containing more than one type of element

Atoms are held together in molecules or compounds by chemical bonds

© McGraw Hill

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Ionic bonds

Formed by the attraction of oppositely charged ions by electrostatic force

Ions form when the atom has a gain or loss of electrons

Na atom loses an electron to become Na+

Cl atom gains an electron to become Cl−

Opposite charges attract so that Na+ and Cl− remain associated as an ionic compound

Electrical attraction of water molecules can disrupt forces holding ions together

© McGraw Hill

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Figure 2.8

© McGraw Hill

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Covalent bonds 1

Form when atoms share 2 or more valence electrons

Results in no net charge, satisfies octet rule, no unpaired electrons

Strength of covalent bond depends on the number of shared electrons

Many biological compounds are composed of more than 2 atoms – may share electrons with 2 or more atoms

© McGraw Hill

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Covalent bonds 2

© McGraw Hill

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Electronegativity

Atom’s affinity for electrons

Differences in electronegativity dictate how electrons are distributed in covalent bonds

Nonpolar covalent bonds = equal sharing of electrons

Polar covalent bonds = unequal sharing of electrons

© McGraw Hill

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Hydrogen bonds

Electropositive hydrogen from one polar molecule is attracted to an electronegative atom that is often oxygen

Attraction produces hydrogen bonds

Each individual bond is weak and transitory

Cumulative effects are enormous

Responsible for many of water’s important physical properties

© McGraw Hill

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Van de Waals Attraction

Weak bond

Non-directional attractive force called Van der Waals forces

Form when two atoms are very close to one another

Antibodies recognize the shape of an invading organism with this bond

© McGraw Hill

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Chemical reactions 1

Chemical reactions involve the formation or breaking of chemical bonds

Atoms shift from one molecule to another without any change in number or identity of atoms

Reactants = original molecules

Products = molecules resulting from reaction

© McGraw Hill

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Chemical reactions 2

Extent of chemical reaction influenced by

Temperature

Concentration of reactants and products

Catalysts

Many reactions are reversible

© McGraw Hill

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Water

Life is inextricably tied to water

Single most outstanding chemical property of water is its ability to form hydrogen bonds

Weak chemical associations that form between the partially negative O atoms and the partially positive H atoms of two water molecules

© McGraw Hill

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Figure

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