Study Guide: Test 1
Study Guide: Test 1 CHEM-111
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This 4 page Study Guide was uploaded by Madelyne Crawford on Saturday September 17, 2016. The Study Guide belongs to CHEM-111 at Campbell University taught by Dr. Kesling in Fall 2016. Since its upload, it has received 43 views. For similar materials see General Chemistry in Chemistry at Campbell University.
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Date Created: 09/17/16
General Chem: Test 1 Study Guide (Chapter 1-3) By Madelyne Crawford Vocabulary Hypothesis- an interpretation of observations; explains scientific findings Law- explains what nature does Theory- explains WHY nature behaves in a certain way Solid- atoms or molecules packed tight together; fixed volume and rigid shape Liquids- pack close, but are free to move; fixed volume, but not a fixed shape Gas- lots of space, compressible; ALWAYS take on shape and volume of containers Crystalline solid- long-range order in terms of atoms and molecules o Table salt, diamonds Amorphous solid- no long-range order o Glass, plastic Pure substance- one component, composition does not vary from sample to sample o Ex: water Mixture- two or more components that can vary (ex: sweet tea) Element vs. compound- cannot be broken down into simpler things (helium) vs. composed of two or more elements in a fixed amount (water) Homogeneous vs. heterogeneous- composition that differs from one part of the mixture to another vs. same composition throughout the mixture Physical vs. chemical property- displayed without going through physical change vs. displayed by going through a chemical change Energy- the capacity to do work and measured in Joules Work- action of a force through a distance Intensive vs. extensive properties- independent of the amount of a substance (density) vs. dependent on the amount of a substance (mass) Precision vs. accuracy- how close measurements are to each other and their reproducibility vs. how close measured value is to the actual value Matter- anything that occupies space and has mass Cathode rays- beams of electrons that come out of the cathode ray tube invented by JJ Thomson Electrostatic forces- compose atoms, which result in attractive and repulsive forces Electric field- area around a charged particle where electrons are Plum-pudding model- JJ Thomson’s idea; negatively charged electrons were tiny particles within a positively charged sphere Atomic number- # of protons in an atom (symbol is Z) Chemical symbol- one or two letter abbreviation (Gold is Au) Isotope- atoms with the same number of protons, but different number of neutrons Natural abundance- the natural amount of isotopes an element has Mass number- the sume of the number of neutrons and protons in an atoms (A) Ions- atoms lose and gain electrons, becoming charged particles Cations- positively charged ions Anions- negatively charged ions Periodic law- when the elements are arranged in order of increasing mass, certain sets of properties recur periodically Metals- lower left and middle of periodic table; good conductors of heat and electricity, malleability, ductility, shiny, lose electrons during chemical reactions Nonmetals- upper right side of periodic table; poor conductors of heat and electricity, gain electrons during reactions Metalloids- along zigzag on periodic table; semiconductors Noble gases- 8A; mostly unreactive Alkaline metals- 1A; all reactive Alkaline earth metals- fairly unreactive Halogens- 7A; very reactive nonmetals Mass spectrometry- method that separates particles in terms of their mass Molar mass- mass on one mole of atoms of an element Ionic bond- metals and nonmetals; TRANSFER of electrons Covalent bond- two or more nonmetals; SHARING of electrons Ionic compound- the result of ionic bonds; lattice of alternating cations and anions Molecule- composed of covalently bonded atoms Empirical formula- relative # of atoms Molecular formula- actual # of atoms Structural formula- uses lines to represent covalent bonds and shows how atoms in a molecule connect or bond to one another Ball and stick model- atoms= balls, bond= sticks; reflects a molecule’s shape Space-filling model- atoms fill up the space between each other to more accurately show the best guess for how a molecule may look if it were a visible site Atomic elements- found in nature with their basic units being single atoms Molecular elements- not usually found in nature with their basic units being single atoms Formal unit- basic unit of an ionic compound; smallest neutral group of ions Binary compound- only have two different elements Oxyanions- anions that have oxygen and another elements Hydrates- have a specific # of water molecules associate with each formal unit Acids- molecular compounds that put out hydrogen ions when dissolved in water; sour taste and ability to dissolve many metals Inorganic nomenclature- having to name a compound without knowing what category it is part of Reactants- substances on left side of chemical reaction Products- right side of chemical reaction Organic compounds- made up of carbon and hydrogen (found in every substance) Alkanes- hydrocarbons containing only single bonds Alkenes- hydrocarbons containing double bonds Alkynes- hydrocarbons containing triple bonds Important People Plato (427-347 BCE) Greek philosopher; thought the best way to learn about reality was through reason instead of the senses Antoine Lavoisier- French chemist; study combustion and how there was no change in total mass of material in container during combustion John Dalton- atomic theory Robert Brown- observed particles in continuous motion; lead to Einstein’s Brownian Motion and the idea of thermal energy Jean Perrin- confirmed Brownian motion and won a Nobel Prize for it (ended doubt that the public had with the presence of particles) Leucippus and Democritus- first to propose the idea of matter being composed of small, indestructible particles Copernicus- marked the beginning of science in the modern world; “On the Revolution of Heavenly Spheres”; idea that the sun was the center of the universe instead of earth Joseph Proust- discovered the composition of compounds; 1797; always in fixed proportions; law of definite proportions JJ Thomson- discovered cathode rays and anodes; invented the cathode ray tube Robert Millikan- oil drop experiment that determined that a charge must be a whole number that is a whole number multiple of the electron’s charge (1909) Ernest Rutherford- nuclear theory; attempted to disprove Thomson James Chadwick- discovered the idea of neutrons Dmitri Mendeleev- developed the first periodic table; elements arranged based on similar properties Theories & Laws Law of conservation of mass- in a chemical reaction, matter can neither be created nor destroyed Law of conservation of energy- energy is neither created nor destroyed Brownian motion- particles encountering thermal energy that keep them in constant motion Law of definite proportions- all samples of a given compound, regardless of their source or how they were prepared, have the same properties of their constituent elements Law of multiple proportions- when two elements form to different compounds, the masses of element B that combine with 1g of element A can be expressed as a ration of small whole numbers Atomic theory o 1. Elements composed of tiny, indestructible particles called atoms o 2. All specific atoms of a certain element have the same mass o 3. Combine in whole number ratios to construct compounds o 4. Not interchangeable (they can only change the way they are bound with other atoms) nuclear theory- o 1. Nucleus contains most of atom’s mass; all positive charge in core o 2. Volume is mostly empty space; negative charged atoms spread throughout o 3. Negative charge in abundance on outside nucleus and protons inside Quizlet Links o Chapter 1: https://quizlet.com/147316410/general-chemistry-chapter- 1-flash-cards/ o Chapter 2: https://quizlet.com/148340884/general-chemistry-chapter- 2-flash-cards/ o Chapter 3: https://quizlet.com/150669212/chem-chapter-3-flash-cards/
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