chem100notes.docx

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Department
Chemistry
Course
CHEM 101
Professor
Beatrice Botch
Semester
Fall

Description
Chemistry 100  Chemistry: The study of matter, transformations, and energy changes  Macroscopic: matter that make up the world and can be seen with the eye  Microscopic: seen under microscope  Submicroscopic: CANNOT be seen under microscope  Molecules: groups of atoms joined in a definite wayforces are covalent bonds o Symbolically represented by molecular formulas o Gives symbol of each atom in the molecule (space-filling and ball and stick)  Compound: made of 2+ elements combined in a definite way  Element: substance made of 1 atom that combine to give molecules  Solids held together by attractive forces between themvibrate in fixed positions (low kinetic energy)  Liquids held together by looser alternative forcesmore freerkinetic energy compared with attraction forces  Gases have no attractive force  Physical properties are characteristic of given substance and can be observed/measured without changing chemical identity (e.g. conductivity, reflectivity, density, melting/boiling temperature)  Physical change takes place without altering chemical identity  Chemical change: conversion of substances (Reactants) to new substance(s) (products) take place  Products have different physical propertiesdifferent appearance (e.g. color change, heat/light) o Physical: mercury is liquid at room temperature o Chemical: methane burns in the air  Matter: can be classified as a pure substance or mixture  Pure substance has same compositioncannot be separated into smaller substances  E.g. elements, compounds, all molecules identical  Mixtures can be separated by physical meansclassified as homo/heterogeneous  Homogeneous: uniform appearance because components mixed evenly o Look like pure substancesaka “solution” (randomly distributed molecules)  Heterogeneous: no uniform appearance because composition distributed unevenlysometimes cannot be seen with the naked eye  Usually opaque or cloudy mixturecan be separated  E.g. areas in container that contain only one type of molecule Elements, atoms, subatomic particles  Different forms of the same element are called allotropes  Atoms are made up of protons, neutrons, and electrons  Subatomic particle are expressed in atomic mass units  Proton and neutron weight 1 AMU, electron 0 AMU  Located in center of atom (nucleus) Isotopes  The number of negatively charged electrons = #positively charged ions  All atoms of same element have same # protons (#protons – atomic #)  Sum of protons and neutrons is mass # (Isotope ID) o Protons = atomic # o Electrons = protons o Mass # = protons and neutrons o Neutrons = mass # - protons o One AMU = mass of one proton or electron o Atomic mass = (fraction 1X atomic mass 1) + (fraction 2 x atomic mass 2)  Weighted average over mass of both isotopes  Periodic table  Elements arranged in order of increasing atomic #  Rows = periods  Columns = groups (common properties)  Elements 1A – 8A = main group  1B-8B: transition elements  Three categories: metals, non-metals, metalloids  Diatomic Molecules: H2, N2, O2, F2, Cl2, Br2, I2  Monatomic gases: He, Ne, Ar, Kr, Xe  Negatively charged ion: ion forms from one atom o E.g. Aluminum atomic # is 13, so 13 protons and 13 electrons o Nucleus has +13 charge, balanced by -13 electrons  Naming monatomic ions from main group elements  Al3+ = aluminum ion  Na+ = sodium ion  Cl- = chloride ion  O2- = oxide ion  Magnitudes  Group 1A metals form ions w/+1 charge  Group 2A metals form ions w/+2 charge  Group 3A metals form ions w/+3 charge  Group 5A metals form ions w/-3 charge  Group 6A metals form ions w/-2 charge  Group 7A metals form ions w/-1 charge  Fe2+ = iron (II) ion  Fe3+ = iron (III) ion  Metals lose electrons and form cations in ionic compounds o Cation name: element name + ion o Cation charge: group #  Nonmetals gain electrons and form anions in ionic compounds o Monatomic anions of n
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