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Essential Chemistry for Biology

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Presentation on theme: "Essential Chemistry for Biology"— Presentation transcript:

1 Essential Chemistry for Biology
CHAPTER 2 Essential Chemistry for Biology

2 Matter: Elements and Compounds
Matter is anything that occupies space and has mass. Matter is found on the Earth in three physical states: Solid Liquid Gas

3 Matter is composed of chemical elements.
Elements are substances that cannot be broken down into other substances. There are 92 naturally occurring elements on Earth. All the elements are listed in the periodic table.

4 Figure 2.2

5 Twenty-five elements are essential to life.
Four of these make up about 96% of the weight of the human body. YOU NEED TO KNOW NHCOPS Trace elements occur in smaller amounts.

6 Figure 2.3

7 Trace elements are essential for life.
An iodine deficiency causes goiter.

8 Elements can combine to form compounds.
These are substances that contain two or more elements in a fixed ratio. Example: NaCl (salt)

9 Each element consists of one kind of atom.
Atoms Each element consists of one kind of atom. An atom is the smallest unit of matter that still retains the properties of an element.

10 Atoms are composed of subatomic particles.
The Structure of Atoms Atoms are composed of subatomic particles. A proton is positively charged. An electron is negatively charged. A neutron is electrically neutral.

11 Elements differ in the number of subatomic particles in their atoms.
The number of protons, the atomic number, determines which element it is. An atom’s mass number is the sum of the number of protons and neutrons. Mass is a measure of the amount of matter in an object.

12 Isotopes are alternate mass forms of an element.
They have the same number of protons and electrons. But they have a different number of neutrons.

13 In radioactive isotopes,
The nucleus decays, giving off particles and energy. Radioactive isotopes have many uses in research and medicine like PET scans Used in dating ancient fossils Uncontrolled exposure to radioactive isotopes can harm living organisms by damaging DNA. Example: the 1999 Chernobyl nuclear accident

14 Electron Arrangement and the Chemical Properties of Atoms
Electrons determine how an atom behaves when it encounters other atoms. Electrons orbit the nucleus of an atom in specific electron shells. The number of electrons in the outermost shell determines the chemical properties of an atom. Magic number is 8 Copyright © 2007 Pearson Education Inc., publishing as Pearson Benjamin Cummings

15 Figure 2.7

16 Chemical Bonding and Molecules
Chemical reactions enable atoms to give up or acquire electrons in order to complete their outer shells. These interactions usually result in atoms staying close together. The atoms are held together by chemical bonds.

17 Ionic Bonds When an atom loses or gains electrons, it becomes electrically charged. Charged atoms are called ions. Ionic bonds are formed between oppositely charged ions. Ionic Bonds

18 Covalent Bonds A covalent bond forms when two atoms share one or more pairs of outer-shell electrons. Covalent Bonds

19 Hydrogen Bonds Water is a compound in which the electrons in its covalent bonds are shared unequally. This causes it to be a polar molecule, one with opposite charges on opposite ends.

20 The polarity of water results in weak electrical attractions between neighboring water molecules.
These interactions are called hydrogen bonds. (Video)

21 Chemical reactions include:
Cells constantly rearrange molecules by breaking existing chemical bonds and forming new ones. Such changes in the chemical composition of matter are called chemical reactions. Chemical reactions include: Reactants, the starting materials Products, the end materials Chemical reactions cannot create or destroy matter, They only rearrange it.

22 Life on Earth began in water and evolved there for 3 billion years.
Water and Life Life on Earth began in water and evolved there for 3 billion years. Modern life still remains tied to water. Your cells are composed of 70%–95% water. The abundance of water is a major reason Earth is habitable.

23 Water’s Life-Supporting Properties
The polarity of water molecules and the hydrogen bonding that results explain most of water’s life-supporting properties: Water’s cohesive nature Water’s ability to moderate temperature Floating ice Versatility of water as a solvent

24 Water molecules stick together as a result of hydrogen bonding.
The Cohesion of Water Water molecules stick together as a result of hydrogen bonding. This is called cohesion. Cohesion is vital for water transport in plants. Adhesion occurs when water sticks to another surface Water Transport

25 Figure 2.12

26 Surface tension is the measure of how difficult it is to stretch or break the surface of a liquid.
Hydrogen bonds give water an unusually high surface tension.

27 How Water Moderates Temperature
Because of hydrogen bonding, water has a strong resistance to temperature change. Water can moderate temperatures. Earth’s giant water supply causes temperatures to stay within limits that permit life. Evaporative cooling removes heat from the Earth and from organisms.

28 The Biological Significance of Ice Floating
When water molecules get cold, they move apart, forming ice. A chunk of ice has fewer molecules than an equal volume of liquid water. The density of ice is lower than liquid water. This is why ice floats.

29 Figure 2.15

30 Since ice floats, ponds, lakes, and even the oceans do not freeze solid.
Marine life could not survive if bodies of water froze solid.

31 Water as the Solvent of Life
A solution is a liquid consisting of two or more substances evenly mixed. The dissolving agent is called the solvent. The dissolved substance is called the solute. When water is the solvent, the result is an aqueous solution.

32 Figure 2.16

33 To describe the acidity of a solution, we use the pH scale.
Acids, Bases, and pH Acid A chemical compound that donates H+ ions to solutions. Base A compound that accepts H+ ions and removes them from solution. To describe the acidity of a solution, we use the pH scale.

34 Figure 2.17

35 Buffers are substances that resist pH change.
They accept H+ ions when they are in excess. They donate H+ ions when they are depleted. Buffering is not foolproof. Example: acid precipitation

36 Figure 2.18

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