Essential Chemistry in Biology: Understanding Matter, Elements, and Atoms

1. CHAPTER 2 Essential Chemistry for Biology

2. Did you know…?

3. Why do we care about Chemistry in Biology? • Take any biological system apart, and you eventually end up at the chemical level.

4. Matter & Elements • Matter:is anything that occupies space and has mass. • Matter is found on the Earth in three physical states: • Solid 2. Liquid 3. Gas • 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.

5. Reading the Period Table Note: You have a periodic table in your text in the appendix at the end on page A-3

6. Elements • Twenty-five elements are essential to life. • Four of these make up about 96% of the weight of the human body. CHON • Trace elements occur in smaller amounts but are still essential for life! • For example iodine deficiency causes goiter. • This is why we add iodine to salt

7. Amounts of Elements in the Body What are the major elements in the air?

8. Elements in the Air

9. Key Element Abbreviations – might be on a QUIZ! Carbon – Hydrogen – Oxygen – Nitrogen – Sulfur – Phosphorus –  Magnesium –  Calcium – Sodium – Potassium – Chlorine – Chloride Ion – Iron – Copper – C H O N S P  Mg  Ca Na K Cl Cl- Fe Cu

10. 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.

11. Compounds • Elements can combine to form compounds. • These are substances that contain two or more elements in a fixed ratio. • Compounds have their own unique properties that are different from those of the contributing elements. +  Na Cl NaCl

12. The Structure of Atoms • Atoms are composed of subatomic particles. • A proton is positively charged. p+ • An electron is negatively charged. e- • A neutron is electrically neutral. n0

13. The Structure of Atoms • Most atoms have protons and neutrons packed tightly into the nucleus. • The nucleus is the atom’s central core. • The electrons orbit the nucleus in a cloud.

14. The Structure of Atoms • Elements differ in the number of subatomic particles in their atoms. • The number of protons, the atomic number, determines which element it is. The number of protons NEVER changes • An atom’s mass number (atomic mass) is the sum of the number of protons and neutrons. • Mass is a measure of the amount of matter in an object without the effect of gravity. This is DIFFERENT than weight!

15. Reading the Periodic Table: Practice • (Mass #) 23 • (Atomic #)11 # of electrons ____ # of protons ____ # of neutrons ____ Na What is the charge of this atom?_____ # of electrons ____ # of protons ____ # of neutrons ____ C (Mass #) 12 (Atomic #) 6 What is the charge of this atom?_____

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

17. Isotopes: Practice (Periodic Table page A-3 in text) • (Mass #) ____ • (Atomic #) ____ # of electrons ____ # of protons ____ # of neutrons ____ 3 H What is the charge of this atom?_____

18. Isotopes: Radioactivity • Radioactive isotopes are unstable: The nucleus decays, giving off particles and energy. • Radioactive isotopes have many uses in research and medicine. • Example: PET scans • Uncontrolled exposure to radioactive isotopes can harm living organisms by damaging DNA. • Example: the 1986 Chernobyl nuclear accident

19. How a PET scan is used: • Alzheimer’s disease is a devastating illness. • It gradually destroys a person’s memory and ability to think. Doctors are now using PET scans in an attempt to diagnose Alzheimer’s disease before the onset of symptoms.

20. Isotopes: Many Applications Who cares? • Dating of geological strata and fossils • Radiation from certain isotopes can be used to treat cancer • Some can kill bacteria that causes food to spoil • Used as labels or tracers to follow the movements of substances within the organisms. Used in research & disease diagnoses. PET scans

21. 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.

22. Figure 2.7

23. 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.

24. 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. • Note: A positive ion has lost an electron NOT gained a proton! Ionic Bonds

26. Atomic Mass: Practice (Periodic Table page A-3 in text) • (Mass #) ? • (Atomic #)? # of electrons ____ # of protons ____ # of neutrons ____ Al+3 What is the charge of this atom?

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

28. 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. • The polarity of water results in weak electrical attractions between neighboring water molecules. • These interactions are called hydrogen bonds.

29. Chemical Reactions • 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.

30. Chemical reactions include: • Reactants: the starting materials • Products: the end materials • Chemical reactions cannot create or destroy matter • They only rearrange it. • This is the Law of Conservation of Matter • Balancing Equations

31. 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.

32. Water’s Life-Supporting Properties Water is Polar Water Structure & Polarity

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

34. 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. • Water Molecules is also attracted to other molecules and this is known as adhesion. Capillary action is demonstrated by this property.

36. 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.

37. How Water Moderates Temperature • Because of hydrogen bonding, water has a strong resistance to temperature change. • Heat and temperature are related, but different. • Heat is the amount of energy associated with the movement of the atoms and molecules in a body of matter. • Temperature measures the intensity of heat. • Water can absorb and store large amounts of heat while only changing a few degrees in temperature.

38. How Water Moderates Temperature • 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.

39. 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. • Did you know that ice can go from a solid to a gas using a process known as sublimation?

40. Why is the density of ice important? • Since ice floats, ponds, lakes, and even the oceans do not freeze solid. • Marine life could not survive if bodies of water froze solid.

41. Water as the Solvent of Life • A solution is a liquid consisting of two or more substances evenly mixed. It is a type of mixture and therefore is NOT chemically combined! • 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.

42. Acids and Bases • Dissociation of water • Breaking apart of the water molecule into two ions of opposite charge (due to strong attraction of oxygen atom of one molecule for H atom of another water molecule) H2O  H+ (hydrogen ion) + OH- (hydroxide ion) H+ +H2O  H3O (hydronium ion)

43. Acids, Bases, and pH • Acid - A chemical compound that donates H+ ions to solutions. In other words an acid is any solution where the number of hydronium ions (H3O) in solutions is greater than the number of hydroxide ions (OH-) • Base - A compound that accepts H+ ions and removes them from solution. In other words a solution is a base if the number of hydroxide ions (OH- ) in solution is greater than the number of hydronium ions • To describe the acidity of a solution, we use the pH scale.

44. pH scale: logarithmic scale for comparing the relative concentrations of hydronium ions and hydroxide ions in a solution ranges from 0 to 14 • the lower the pH the stronger the acid • the higher the pH the stronger the base • pH 7.0 is neutral

45. Buffers • Buffers are substances that resist pH change. Typically they are weak acids and bases that neutralize excess acids and bases in solution • They accept H+ ions when they are in excess. • They donate H+ ions when they are depleted. • Control of Buffers is CRITICAL for living things • Most enzymes in your body only function within a narrow pH range • Buffering is not foolproof. • Example: acid precipitation 

46. Chemical Equations Part 2 How to read and balance chemical equations

47. Chemical Formulas • Chemical formulas show the elements & number of atoms of each element in a molecule • In small letters to the right of the elements there is a number that indicates how many atoms of that element are in the molecule. This number is called the subscript. If there is no number assume the number is “1” • H2SO4 • How many atoms of each element is present? • Elements • Hydrogen: 2 atoms • Sulfur: 1 atom • Oxygen: 4 atoms • 7 atoms total Subscript

48. Coefficients • A formula may begin with a number. • If there is no number, then “1” is understood to be in front of the formula. • This number is called the coefficient. • The coefficient represents the number of molecules of that compound or atom needed in the reaction. • For example: • 2H2SO4 • The 2 indicates that there are 2 molecules of Sulfuric Acid Coefficient

49. Coefficients • 2H2SO4 – 2 molecules of Sulfuric Acid • A coefficient is distributed to ALL elements in a compound 2 – H2 (multiply 2x2) (for a total of 4 H atoms) 2 – S (multiply 1(assume S has 1) x2) (for a total of 2 S atoms) 2 – O4(multiply 2x4) (for a total of 8 O atoms)

50. WHAT DOES THIS HAVE TO DO WITH BIOLOGY!!?? grrrrrr