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Biological Sciences
Heinz- Bernhard Kraatz

CHEMISTRY: A MOLECULAR APPROACH 1 MATTER, MEASUREMENT, AND PROBLEM SOLVING 1.1 Atoms & Molecules  A carbon monoxide molecule contains a carbon atom and an oxygen atom held together by a chemical bond  Atoms are the submicroscopic particles that constitute the fundamental building blocks of ordinary matter  Molecules are two or more atoms joined in a specific geometrical arrangement  To understand substances around us, we must understand the atoms and molecules that compose them  Chemistry is the science that seeks to understand the behaviour of matter by studying the behaviour of atoms and molecules 1.2 The Scientific Approach to Knowledge  Scientific knowledge is based on observation and experiment  A hypothesis is a tentative interpretation or explanation of the observations  Experiments are highly controlled procedures designed to generate observations  Scientific Law is a brief statement that summarizes past observations and predicts future ones  The Law of Conservation of Mass states, “In a chemical reaction, matter is neither created nor destroyed.”  A law summarizes a series of related observations, while a theory gives the underlying reasons for them 1.3 The Classification of Matter  Matter is anything that occupies space and has mass  A specific instance of matter, such as air, water, or sand is called a substance  Matter is classified according to state ( physical form) and composition  Solid matters maybe crystalline (atoms or molecules are patterns with long range, repeating order) or amphorous (atoms and molecules do not have any long-range order)  A pure substance is made up of only one component and its composition is variant  A mixture is composed of two or more components in varying proportions  Pure substances can be split into elements and compounds  Mixture can be split into heterogeneous mixtures (composition varies throughout) and homogenous mixtures (same composition throughout)  Mixtures can be separated by decanting, distillation, or filtration 1.4 Physical & Chemical Changes and Physical & Chemical Properties  Physical Changes – changes in state/appearance, not composition  Chemical Changes – changes that alter the composition of matter  Physical Property – a property displayed without changing its composition (odor, taste, colour, appearance, melting point, boiling point, and density)  Chemical Property – a property displayed only by changed composition via a chemical change (corrosiveness, flammability, toxicity) 1.5 Energy: A Fundamental Part of Physical and Chemical Change  Physical and chemical changes are often accompanied by energy changes o To understand chemistry we must understand energy changes and energy flow  Energy is the capacity to do work  Work is the action of a force through a distance  The total energy of an object is the sum of its kinetic energy, the energy associated with its motion is potential energy, and the energy associated with the temperature of an object is thermal energy o Thermal energy is a type of kinetic energy  Law of Conservation of Energy - Energy is neither created nor destroyed  Systems with high potential energy tend to change in a way that lowers their potential energy o Therefore systems with high potential energy tend to be unstable  Chemical potential energy arises primarily from electrostatic forces between the electrically charged particles (protons and electrons) that compose atoms and molecules 1.6 The Units of Measurement  A meter is the distance light travels through a vacuum in a certain period of time o 1 yard = 36 in o 1 m = 39.37 in  The mass of an object is the measure of the quantity of matter within it, while the weight of an object is the measure of gravitational pull on its matter o 1 kg = 2.205 lb o 1 kg = 1000 g  A second is the duration of 9,192,631,770 periods of radiation emitted from a certain transition in a cesium-133 atom  Temperature is the measure of average kinetic energy and determines the direction of thermal energy transfer (heat) o 212 ° F = 100 °C = 373 K (Water boils) o 32 ° F = 0 °C = 273 K (Water boils) o -459 ° F = -273 °C = 0 K (Water boils)  Density is an example of an intensive property, which means that it is independent of the amount of the substance  1 gallon = 3.785 L 1.7 The Reliability of Measurement  Scientific measures are reported so that every digit is certain except the last, which is estimated 1.8 Solving Chemical Problems  No one succeed in life without the ability to solve problems 2 ATOMS AND ELEMENTS  The word atom comes from the Greek atomos meaning “indivisible.” 2.1 Imaging and Moving Individual Atoms  Scanning Tunneling Microscopy (STM) is technique that can image and even move individual atoms and molecules o It works by moving an extremely sharp electrode (an electric conductor) over a surface and measuring the resulting tunnel current, the electrical current that flows between the tip of an electrode and the surface even though the two are not in physical contact  An atom is the smallest identifiable unit of an element 2.2 Early Ideas about the Building Blocks of Matter  The first people to propose matter was composed of small, indestructible particles were Leucippus and his student Democritus o They theorized that matter was composed of indivisible particles they named atomos  Plato and Aristotle held that matter had no small parts and that different substances were composed of various proportions of fire, air, earth, and water  In the 1800s John Dalton offered convincing evidence that supported the idea of the atom 2.3 Modern Atomic Theory and the Laws That Led to It  The Law of the Conservation of Mass (Antoine Lavoisier 1789) – In a chemical reaction, matter is neither created nor destroyed  The Law of Definite Proportions (Joseph Proust 1797) – All samples of a given compound, regardless of their source or how they were prepared, have the same proportions as their constituent elements  Law of Multiple Proportions (John Dalton 1804) – When two elements (call them A and B) form two different compounds, the masses of element B that combine with 1 g of element A can be expressed as a ratio of small whole numbers  All these laws were supported by John Dalton’s atomic theory in 1808, which includes these concepts: 1) Each element is composed of tiny, indestructible particles called atoms 2) All atoms of a given element have the same mass and other properties that distinguish them from the atoms of other elements 3) Atoms combine in simple, whole-number ratios to form compounds 4) Atoms of one element cannot change into atoms of another element. In a chemical reaction, atoms only change the way they are bound together with other atoms. 2.4 The Discovery of the Electron  In the late 1800s, J. J. Thompson discovered the electron using a cathode ray tube  In 1909, Millikan determined the charge of the electron using the Oil Drop Experiment o The charge of an electron was determined to be  The magnitude of the charge of an electron is important because it determines how strongly an atom holds its electrons 2.5 The Structure of the Atom  J. J. Thompson proposed that the negatively charged electrons were small particles held within a positively charged sphere (plum-pudding model aka blueberry muffin where the blueberries are electrons and the muffin is the positively charged sphere)  Radioactivity is the emission of small energetic particles from the core of certain unstable atoms  In 1909, Ernest Rutherford directed positively charged particles at an ultrathin sheet of gold foil o A majority of the particles were deflected, but approximately 1 in 20,000 bounced back o He realized that the mass and positive charge of an atom must be concentrated in a space much smaller than the size of the atom itself o Rutherford proposed the nuclear theory of the atom, with three basic parts:  Most of the atom’s mass and all of its positive charge are contained in a small core called the nucleus  Most of the atom’s volume is empty space where tiny, negatively charged electrons are dispersed  There are equal amounts of negatively charged particles outside the nucleus as there are positively charged particles inside the nucleus. This makes the atom electrically neutral.  Scientist realized this Rutherford’s model was incomplete due to varied mass from one atom to another. o Rutherford and Chadwick discovered the unaccounted weight was due to neutrons  The dense nucleus accounts for 99.9% of the mass of an atom, but occupies very little of its volume  Matter, at its core, is much less unified than it appears  Matter appears solid because the variation in its density is on such a small scale that our eyes cannot see it 2.6 Subatomic Particles: Protons, Neutrons, and Electrons in Atoms  A more common unit to express the masses of the neutron and proton is the atomic mass unit (amu), defined as 1/12 the mass of a carbon atom containing six protons and six neutrons. o Using this unit, the mass of a proton/neutron is approximately 1 amu  The charge of the proton and the electron is opposite in sign but equal in magnitude  The number of particles is what makes the difference between atoms of one element compared to the atoms of another element o The most important identity of an atom is the number of protons in its nucleus  The atomic number is given the symbol Z  The elements are arranged so that those with similar properties are in the same column  Each element is represented with a unique chemical symbol  Atoms with the same number of protons but a different number of neutrons are called isotopes o An isotope with more neutrons has a greater mass  Mass number (A) = # of protons (p) + # of neutrons (n)  Cations are positively charged ions, negatively charged ions are called anions 2.7 Finding Patterns: The Periodic Table and the Periodic Law  Mendeleev was an important figure in the creation of the modern periodic table  He determined that when the elements are arranged in order of increasing mass, certain sets of properties recur periodically (periodic law)  In the modern periodic table, elements are in order of increasing atomic radius rather than increasing atomic mass  Elements can be classified as metals, non-metals, and metalloids o A main group metal tends to lose electrons forming a cation o A main group non-metal tends to gain electrons forming an anion 2.8 Atomic Mass: The Average Mass of an Element’s Atoms  Atomic mass = fraction of isotope 1 mass of isotope 1 + fraction of isotope 2 mass of isotope 2 + fraction of isotope 3 mass of isotope 3 + …  The masses of atoms and percent abundances of isotopes of elements are measured using mass spectrometry 2.9 Molar Mas: Counting Atoms by Weighing Them  The value of a mole is equal to the number of atoms in exactly 12 grams of pure carbon-12 (12 g C = 1 mol C atoms = C)  An element’s molar mass in grams/mole is numerically equal to the element’s atomic mass in atomic mass unit Do we need to know the laws? Do we need to know JJ Thompson and Millikan experiments? Cations and anions? Do we need to know mass spectrometry? 3 MOLECULES, COMPOUNDS AND CHEMICAL EQUATIONS 3.1 Hydrogen, Oxygen, and Water  Hydrogen and oxygen have extremely low boiling points (-253 C and -183 C)  The properties of compounds are often very different from the properties of the elements that compose them  Free atoms are rare on earth  In a compound, elements combine in fixed definite proportions while in mixtures, elements combine in any proportions whatsoever 3.2 Chemical Bonds  Compounds are composed of atoms held together by chemical bonds, which are the result of interactions between charged particles that compose the atom  Ionic Bond – between metals and non-metals, involve the transfer of electrons from one atom to another o A metal transfers its electrons to the non-metal o The metal becomes a cation and the non-metal becomes an anion, causing attraction between the two elements through electrostatic forces  Covalent Bond – between two or more non-metals, involve the sharing of electrons between two atoms o The shared electrons interact with the nuclei of both atoms, lowering the potential energy (more stability) of the system through electrostatic interactions 3.3 Representing Compounds: Chemical Formulas and Molecular Models  The empirical formula gives the relative number of atoms of each element in a compound ( )  The molecular formula gives the actual number of atoms of each element in a molecule of a compound ( )  A structural formula uses lines to represent covalent bonds and shows how atoms in a molecule are connected/bonded to each other, and can be used to show a molecules geometry 3.4 An Atomic-Level View of Elements and Compounds  Atomic elements are those that exist in nature with single atoms as their basic units  Molecular elements normally exist in nature as molecules (ex. diatomi
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