Bettelheim, Brown, Campbell and Farrell Chapter 2 Atoms Bettelheim, Brown, Campbell and Farrell Chapter 2

Classification of Matter

Classification of Matter Element: Pure substance made up of “identical” atoms 116 known elements 88 occur in nature; others man-made Represented by one or two letter symbols First letter CAPITALIZED; second letter small

Elements found in Nature Monatomic elements: Exist as single atoms Diatomic elements: Occur as diatomic molecules (pairs of atoms) H2, N2, O2, F2, Cl2, Br2, and I2 Polyatomic elements: Have three or more atoms per molecule O3, P4, S8, diamond

Classification of Matter Compound: Pure substance made up of two or more elements in a fixed ratio by mass Formula of a compound: Gives the ratio of each element (uses atomic symbols). Examples: NaCl H2O

Classification of Matter Mixture: Combination of two or more pure substances Substances may be present in any mass ratio Each substance has a different set of physical properties Each substance keeps its own properties and identity Can separate mixture into the individual substances by using the physical properties of the individual substances in the mixture

Types of Mixtures Heterogeneous mixture: Substances are not evenly distributed throughout Homogeneous mixture (Solution): Substances are evenly distributed throughout.

Dalton’s Atomic Theory - 1805 All matter is composed of very tiny particles (atoms) All atoms of the same element are the same (same chemical properties) Compounds are formed by the chemical combination of two or more different kinds of atoms A molecule is a tightly bound combination of two or more atoms that acts as a unit

Evidence for Dalton’s Theory Law of Conservation of Mass Mass can be neither created or destroyed Law of Constant Composition Compounds have a definite composition by mass

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Subatomic Particles The unit of mass is the atomic mass unit (amu) one amu is defined as one-twelfth the mass of an atom of carbon with 6 protons and 6 neutrons in its nucleus 1 amu = 1.6605 x 10-24 g

A typical atom

Mass and Atomic Numbers Mass number: Sum of the number of protons plus neutrons in the nucleus of an atom Atomic number: Number of protons in the nucleus of an atom Notation for single atom or isotope

Isotopes Isotopes: atoms with the same number of protons but a different number of neutrons Most elements found as mixtures of isotopes Atomic weight: weighted average of masses of isotopes of an element

Classification of Elements Metals Solids (except Hg), shiny, conduct electricity, ductile, and malleable Tend to give up electrons to form positive ions Nonmetals On right side of Periodic Table (except H) Brittle, dull, poor conductors of electricity Tend to accept electrons to form negative ions

Classification of Elements Metalloids (also called semi-metals) B, Si, Ge, As, Sb, Te Have some properties of metals and some of nonmetals Silicon is a semiconductor Does not conduct electricity at low voltages, but becomes a conductor at higher voltages

Classification of Elements

Classification of Elements, part 2 Much info known about elements by mid 1800s Tried to organize elements in logical way Dmitri Mendeleev’s pattern worked best Noted that certain properties tended to recur periodically Took elements in order of increasing mass Placed elements with similar properties in same column

Modern Periodic Table Elements arranged in order of increasing atomic number (# of protons) Elements with repeating properties are in the same group or family (column or vertical row) Elements in same period (horizontal row) change properties as you go across the period

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Periodic Properties Mendeleev looked at both chemical and physical properties

Trend within Same Group The alkali metals, Group 1A elements Melting and boiling points decrease as you go down table

Group 1A: Alkali Metals React with halogens to form compounds such as NaCl Form +1 ions Very reactive

Examples of Periodicity Fluorine, chlorine, bromine, and iodine fall in the same column

Examples of Trends in Periodicity The halogens, Group 7A elements Melting and boiling points increase as you go down table

Group 7A: Halogens Halogens exist as diatomic molecules, such as Cl2, F2, etc Halogens react with group I metals compounds such as NaCl Form -1 ions Very reactive

Examples of Periodicity The noble gases, Group 8A elements Melting and boiling points increase as you go down table

Group 8A: Noble Gases Do not react with other elements Do not form ions

Why do elements in group have similar properties? Outermost Electron configuration is the same for all elements in a group Electron configuration: the arrangement of electrons outside the nucleus

Electron Configuration Electrons are distributed in shells about the nucleus

Electron Configuration Each shell (principal energy level) has a different maximum number of electrons it can hold

Electron Configuration Shells (principal energy levels) are subdivided into orbitals

Electron Configuration Different kinds of orbitals have definite (and different) shapes and orientations in space

Electron Configuration Each orbital has a different energy Lowest energy fills first

Electron Configuration: Building Atoms Rule 1: Lowest energy orbitals fill first The first three energy level orbitals fill in the order 1s, 2s, 2p, 3s, and 3p Rule 2: Each orbital can hold a maximum of two electrons (with opposite spins) Rule 3: Orbitals of equal energy each add one electron first, then add second electron to fill them completely.

Order of Filling Orbitals 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s Don’t just fill up one shell and then the next shell completely Lowest energy fill first

Types of orbitals per shell

Orbital # orbitals/shell # electrons/ orbital type s 1 2 p 3 6 d 5 10 f 7 14

Electron Configuration Orbital box diagrams a box represents an orbital an arrow represents an electron a pair of arrows with heads in opposite directions represents a pair of electrons with paired spins Example: carbon (atomic number 6)

Electron Configuration Notation (Compact shorthand) 3p4 3 is principal energy level (shell) p is type of orbital (subshell) Superscript 4 show that there are four electrons in 3p orbitals

Electron Configuration Noble gas notation the symbol of the noble gas immediately preceding the particular atom indicates the electron configuration of all filled shells Example: carbon (atomic number 6)

Write the electron configuration for Vanadium (# 23)

Shorthand way of remembering 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f

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Electron Configuration Lewis dot structures for the Group 1A (alkali) metals

Lewis Dot Structures Atomic symbol represents nucleus and core electrons Valence electrons shown as dots around symbol

Ionization Energy Ionization energy: the energy required to remove the most loosely held electron from an atom in the gaseous state example: when lithium loses one electron, it becomes a lithium ion; it still has three protons in its nucleus, but now only two electrons outside the nucleus

Ionization Energy Ionization energy is a periodic property

Ionization Energy Ionization energy is a periodic property Increases as you move from left to right Increases as you move up the table

Electron Configuration Valence shell: the outermost incomplete shell Valence electron: an electron in the valence shell Core electron: Electron inside the outermost shell

Electrons in Energy Levels The energy of electrons in an atom is quantized An electron can have only certain allowed energies Ground state: the electron configuration of lowest energy Excited state: Electron has more than lowest possible energy