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  2. Definition of Concepts Matter and Energy

  3. Matter • Is anything that occupies space and has mass • The mass of an object, which is equal to the actual amount of matter in the object, remains constant wherever the object is • In contrast, weight varies with gravity • Remains constant regardless of gravity • Weight does not

  4. States of Matter • Matter exists in one of three states: • Solid • Liquid • gas

  5. ENERGY • Has no mass and does not take up space • Compared with matter, energy is less tangible • Measured by only its effect on matter • Is the capacity to do work, or to put matter into motion

  6. ENERGY • Exists in two forms, or work capacities, each transformable to the other: • Kinetic energy: energy of motion • Energy in action • Potential energy: stored energy • Inactive energy that has the potential, or capability, to do work but is not presently doing so • Matter is the substance, and energy is the mover of the substance

  7. ENERGY • Forms of energy: • Chemical: energy stored in chemical bonds • Potential energy in the foods you eat is eventually converted into the kinetic energy of movement • Food fuels cannot be used to energize body activities directly • Some of the food energy is captured temporarily in the bonds of a chemical called adenosine triphosphate (ATP) • Electrical: results from the movement of charged particles • Electrical currents are generated when charged particles called ions move along or across cell membranes • Nervous system uses electrical currents, called nerve impulses, to transmit messages from one part of the body to another • Mechanical: energy directly involved with moving matter • Walking, running, movement of arms, etc. • Radiant (electromagnetic): energy that travels in waves • Light energy stimulates the retina of the eye • Ultraviolet waves cause sunburn, but they also stimulate our body to make vitamin D • Easily converted from one form to another


  9. BASIC TERMS • Elements are unique substances that cannot be broken down into simpler substances by ordinary chemical means • Four elements: carbon, hydrogen, oxygen, and nitrogen make up roughly 96% of body weight • Atoms are the smallest particles of an element that retain the characteristics of that element • Every element’s atoms differ from those of all other elements and give the element its unique physical (color, texture, boiling point, freezing point) and chemical properties (the way atoms interact with other atoms: bonding behavior) • Elements are designated by a one- or two- letter abbreviation called the atomic symbol

  10. ATOMIC STRUCTURE • Atom: Greek for indivisible • Each atom has a central nucleus with tightly packed protons and neutrons • Protons (p+) have a positive charge and a mass of 1 atomic mass unit (amu) • Neutrons (n0) do not have a charge but have a mass of 1 atomic mass unit (amu) • Thus, the nucleus is positively charged overall • Accounts for nearly the entire mass (99.9%) of the atom • Electrons (e-) are found moving around the nucleus, have a negative charge, and are considered massless (0 amu)????? • 1/2000 the mass of a proton


  12. ATOMIC STRUCTURE • All atoms are electrically neutral because the number of electrons in an atom is equal to the number of protons (the + and – charges cancel the effect of each other) • For any atom the number of protons and electrons is always equal

  13. ATOMIC STRUCTURE • Planetary model (a): is a simplified (outdated), two-dimensional model of atomic structure • It depicts electrons moving around the nucleus in fixed, generally circular orbits • BUT, we can never determine the exact location of electrons at a particular time because they jump around following unknown trajectories


  15. ATOMIC STRUCTURE • Orbital model (b): is a more accurate three dimensional model talking about orbital regions instead of set orbital patterns • Instead of speaking of specific orbits, chemists talk about orbitals—regions around the nucleus in which a given electron pair is likely to be found most of the time • More useful for predicting the chemical behavior of atoms • Depicts probable regionsof greatest density by denser shading (this haze is called the electron cloud)


  17. IDENTIFYING ELEMENTS • Elements are identified based on their number of protons, neutrons, and electrons • All we really need to know to identify a particular element are its atomic number, mass number, and atomic weight


  19. ATOMIC NUMBER • Is equal to the number of protons in the nucleus of any atom • Written as a subscript to the left of its atomic symbol • Examples: • Hydrogen with one proton, has an atomic number of 1 (1H) • Helium with two protons, has an atomic number of 2 (2He) • Since the number of protons is equal to the number of electrons, the atomic number indirectly tells us the number of electrons • This is important information, because electrons determine the chemical activity of atoms

  20. Mass Number and Isotopes • Mass number of an element is equal to the number of protons plus the number of neutrons • The electron is considered massless and is ignored in calculating the mass number • Examples: • Hydrogen has only one proton in its nucleus, so its atomic and mass numbers are the same: 1 • Helium, with two protons and two neutrons, has a mass number of 4 • Mass number is usually indicated by a superscript to the left of the atomic symbol • Thus, helium is: 42He • This simple notation allows us to deduce the total number and kinds of subatomic particles in any atom because it indicates the number of protons (the atomic number), the number of electrons (equal to the atomic number), and the number of neutrons (mass number minus atomic number)

  21. Mass Number and Isotopes • Nearly all known elements have two or more structural variations called isotopes • They have the same number of protons and electrons of all other atoms of the element but differ in the number of neutrons in the atom • Examples: • Hydrogen has a mass number of 1: 1H • Some hydrogen atoms have a mass of 2 or 3 amu, which means that they have one proton and, respectively, one or two neutrons: 2H or 3H


  23. Isotopes • Carbon has several isotopic forms: • The most abundant of these are: 12C, 13C, and 14C • Each of the carbon isotopes has six protons (otherwise it would not be carbon), but 12C has six neutrons, 13C has seven neutrons, and 14C has eight neutrons • Isotopes are also written with the mass number following the symbol: C-14

  24. ATOMIC WEIGHT • Also referred to as ATOMIC MASS • Is an average of the relative masses of all isotopes of an element, taking into account their relative abundance (proportions) in nature • Example: • Atomic mass of hydrogen is 1.008 • Reveals that its lightest isotope (1H) is present in much greater amounts in our world than its 2H or 3H forms

  25. RADIOISOTOPES • The heavier isotopes of many elements are unstable and spontaneously decompose into more stable forms • The process of atomic decay is called radioactivity, and isotopes that exhibit this behavior are called radioisotopes • The disintegration of a radioactive nucleus may be compared to a tiny explosion • It occurs when subatomic alpha (packets of 2p + 2n) particles, beta (electronlike negative particles) particles, or gamma (electromagnetic energy) rays are ejected from the atomic nucleus • Why this happens is complex, and you only need to know that the dense nuclear particles are compressed of even smaller particles called quarks that associate in one way to form protons and in another way to form neutrons • Apparently, the “glue” that holds these nuclear particles together is weaker in the heavier isotopes • When disintegration occurs, the element may transform to a different element

  26. RADIOISOTOPES • Radioisotopes gradually lose their radioactive • Time required for a radioactive isotope to lose one-half of its radioactivity is called the half-life (varies from hours to thousands of years)


  28. MOLECULES AND COMPOUNDS • A combination of two or more atoms is called a molecule • If two or more atoms of the same element combine it is called a molecule of that element • H2,, O2 , S8 • If two or more atoms of different elements combine it is called a molecule of a compound • H2O, CH4 • Just as an atom is the smallest particle of an element that still exhibits the properties of the element, a molecule is the smallest particle of a compound that still displays the specific characteristics of the compound • Important concept: • Because the properties of compounds are usually very different from those of the atoms they contain

  29. MIXTURES • Substances made of two or more components mixed physically • Although most matter in nature exists in the form of mixtures, there are only three basic types: • Solutions • Colloids • suspensions

  30. Solutions • Homogeneous mixtures of compounds that may be gases, liquids, or solids • Examples: • Air: mixture of gases • Seawater: mixture of salts, which are solid, and water • The substance present in the greatest amounts is called the solvent (does the dissolving) • Usually liquids • Water is the universal solvent • Substances present in smaller amounts are called solutes (is dissolved) • Most solutions in the body are true solutions containing gases, liquids, or solids dissolved in water • True solutions are usually transparent • Examples: • Saline solution: NaCl and water • Glucose and water • Solutes of a true solution are minute, usually in the form of individual atoms and molecules • Consequently, they are not visible to the naked eye, do not settle out, and do not scatter light • If a beam of light is passed through a true solution, you will not see the path of light

  31. Concentration of Solutions • Solutions may be described by their concentrations, which may be indicated in various ways: • Percent (parts per 100 parts) of the solute in the solution • Always refers to the solute percentage, and unless otherwise noted, water is assumed to be the solvent • Molarity (moles per liter): • Indicated by M • Mole of any element or compound is equal to its atomic weight or molecular weight (sum of the atomic weights) weighed out in grams

  32. Concentration of SolutionsMolarity • Glucose is C6H12O6, which indicates that it has 6 carbon atoms, 12 hydrogen atoms, and 6 oxygen atoms • The molecular weight of glucose using the periodic table (chart) is calculated as follows: • Atom Number Atomic Total • of Weight Atomic • Atoms Weight • C 6 X 12.011 = 72.066 • H 12 X 1.008 = 12.096 • O 6 X 15.999 = 95.994 • 180.156

  33. Concentration of SolutionsMolarity • To make a one-molar solution of glucose, you would weigh out 180.156 grams (g), called a gram molecular weight, of glucose and add enough water to make 1 liter (L) of solution • Thus, a one-molar solution (1.0 M) of a chemical substance is one gram molecular weight of the substance (or one gram atomic weight in the case of elemental substances) in 1 L (1000 ml) of solution

  34. Concentration of SolutionsMolarity • The beauty of using the mole as the basis of preparing solutions is its precision: • One mole of any substance contains exactly the same number of solute particles, that is, 6.02 X 1023 (Avogadro’s number) • So whether you weigh out 1 mole of glucose (180 g) or 1 mole of water (18 g) or 1 mole of methane (16 g), in each case you will have 6.02 X 1023 molecules of that substance

  35. Colloids • Colloids (emulsions) are heterogeneous mixtures that often appear translucent or milky • Although, the solute particles are larger than those in true solutions, they still do not settle • However, they do scatter light, and so the path of a light beam shining through a colloidal mixture is visible

  36. Colloids • Have many unique properties, including the ability of some to undergo sol-gel transformation, that is, to change reversibly from a fluid (sol) state to a more solid (gel) state • Jell-O, or any gelatin product, is a familiar example of a nonliving colloid that changes from a sol to a gel when refrigerated (and that will liquefy again if placed in the sun) • Cytosol, the semifluid material in living cells, is also a colloid, and its sol-gel changes underlie many important cell activities, such as cell division

  37. Suspensions • Suspensions are heterogeneous mixtures with large, often visible solutes that tend to settle out • Examples: • Mixture of sand and water • Blood: living blood cells are suspended in the fluid portion of blood (blood plasma)

  38. DISTINGUISHING MIXTURES AND COMPOUNDS • 1.The main difference between mixtures and compounds is thatno chemical bonding occurs between molecules of a mixture • Properties of atoms and molecules are not changed when they become part of a mixture • They are ONLY physically intermixed • 2. Mixtures can be separated into their chemical components by physical means (straining, filtering, evaporation, etc.); separation of compounds is done by chemical means (breaking bonds) • 3. Some mixtures are homogeneous, while others are heterogeneous: • Homogenous means that a sample taken from any part of the substance has exactly the same composition (in terms of the atoms or molecules it contains) as any other sample • A bar of 100% pure (elemental) iron is homogeneous, as are all compounds • Heterogeneous substances vary in their makeup from place to place • Iron ore is a heterogeneous mixture that contains iron and many other elements

  39. CHEMICAL BONDS • A chemical bond is an energy relationship between the electrons of the reacting atoms • NOT a physical structure

  40. Role of Electrons in Chemical Bonding • Electrons occupy regions of space called electron shells that surround the nucleus in layers • The atoms known so far can have electrons in seven shells (numbered 1 to 7 from the nucleus outward) • But, the actual number of electron shells occupied in a given atom depends on the number of electrons that atom has • Each electron shell contains one or more orbitals • Each electron shell represents a different energy level (think of electrons as particles with a certain amount of potential energy) • Electron shell and energy level are used interchangeable • Each electron shell represents a different energy level • Each electron shell holds a specific number of electrons, and shells tend to fill consecutively from the closest to the nucleus to the furthest away • The octet rule, or rule of eights, states that except for the first energy shell (stable with two electrons), atoms are stable with eight electrons in their outermost (valence) shell

  41. Role of Electrons in Chemical Bonding • The amount of potential energy an electron has depends on the energy level it occupies, because the attraction between the positively charged nucleus and negatively charged electrons is greatest closest to the nucleus and falls off with increasing distance • This statement explains why electrons farthest from the nucleus: • 1. Have the greatest potential energy (it takes more energy to overcome the nuclear attraction and reach the more distant energy levels) • 2. Are most likely to interact chemically with other atoms (they are the least tightly held by their own atomic nucleus and the most easily influenced by other atoms and molecules

  42. Role of Electrons in Chemical Bonding • Each electron shell can hold a specific number of electrons: • Shell 1: shell immediately surrounding the nucleus • Accommodates only 2 electrons • Shell 2: holds a maximum of 8 • Shell 3: holds a maximum of 18 • Subsequent shells hold larger and larger numbers of electrons • Shells tend to be filled consecutively (from Shell 1 outward)

  43. Role of Electrons in Chemical Bonding • When considering bonding behavior, the only electrons that are important are those in the atom’s outermost energy level • Inner electrons usually do not take part in bonding because they are more tightly held by the atomic nucleus • Before an atom reacts it is electrically stable (same number of protons and electrons) BUT it might not be chemically stable • Chemical stability depends on the outer energy level being filled



  46. Role of Electrons in Chemical Bonding • In atoms that have more than 20 electrons, the energy levels beyond shell 2 can contain more than eight electrons • However, the number of electrons that can participate in bonding is still limited to a total of eight • The term valence shell is used specially to indicate an atom’s outermost energy level or that portion of it containing the electrons that are chemically reactive • Hence, the key to chemical reactivity is the octet rule, or rule of eights • Except for Shell 1, which is full when it has two electrons, atoms tend to interact in such a way that they have eight electrons in their valence shell



  49. Types of Chemical Bonding • Three major types of chemical bonds: • Ionic • Covalent • Hydrogen

  50. Ionic Bonds • Atoms are electrically neutral but might not be chemically stable: • Electrons can be transferred from one atom to another, and when this happens, the precise balance of + and – charges is lost and charged particles called ions are formed • Ionic bonds are chemical bonds that form between two atoms that transfer one or more electrons from one atom to the other • Ions are charged particles • An anion is an electron acceptor carrying a net negative charge due to the extra electron (gains electrons) • A cation is an electron donor carrying a net positive charge due to the loss of an electron (it might help you to think of the “t” in “cation” as a + sign) • Because opposite charges attract, these ions tend to stay close together, resulting in an ionic bond