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Introduction & Atoms, Molecules, and Ions

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1 Introduction & Atoms, Molecules, and Ions
Chapter 1-2: Introduction & Atoms, Molecules, and Ions Chemistry science study matter, its properties, and its behavior. Matter anything that has mass and takes up space. Prof. Fakhr M. Abu-Awwad Islamic University - Gaza

2 Atoms are the building blocks of matter.
Each element is made of the same kind of atom. A compound is made of two or more different kinds of elements.

3 States of Matter

4 Classification of Matter
And Measurement

5 Properties and Changes of Matter
Types of Properties 1 Physical Properties… Can be observed without changing a substance into another substance. Boiling point, density, mass, volume, etc. Chemical Properties… Can only be observed when a substance is changed into another substance. Flammability, corrosiveness, reactivity with acid, etc.

6 Physical Changes Chemical Changes Types of Changes
These are changes in matter that do not change the composition of a substance. Changes of state, temperature, volume, etc. Chemical Changes Chemical changes result in new substances. Combustion, oxidation, decomposition, etc.

7 Chemical Reactions In the course of a chemical reaction, the reacting substances are converted to new substances.

8 Units of Measurement SI Units

9 Metric System Prefixes convert the base units into units that are appropriate for the item being measured.

10 Volume The most commonly used metric units for volume are the liter (L) and the milliliter (mL). A liter is a cube 1 decimeter (dm) long on each side. A milliliter is a cube 1 centimeter (cm) long on each side.

11 Temperature temperature is a measure of the average kinetic energy of the particles in a sample.

12 the Celsius and Kelvin scales are most often used.
The Celsius scale is based on the properties of water. 0 C is the freezing point of water. 100 C is the boiling point of water. The kelvin is the SI unit of temperature. It is based on the properties of gases. There are no negative Kelvin temperatures. K = C F = 9/5(C) + 32 C = 5/9(F − 32)

13 Chapter 2 Atoms, Molecules, and Ions

14 Atomic Theory of Matter
The theory that atoms are the fundamental building blocks of matter reemerged in the early nineteenth century, championed by John Dalton. Dalton's Postulates - Each element is composed of extremely small particles called atoms. - All atoms of a given element are identical to one another in mass and other properties, but the atoms of one element are different from the atoms of all other elements.

15 Law of Conservation of Mass
- Atoms of an element are not changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions. Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms. Law of Conservation of Mass The total mass of substances present at the end of a chemical process is the same as the mass of substances present before the process took place.

16 The Electron Streams of negatively charged particles were found to emanate from cathode tubes, causing fluorescence. J. J. Thomson is credited with their discovery (1897). Thomson measured the charge/mass ratio of the electron to be 1.76  108 coulombs/gram (C/g). Robert Millikan (University of Chicago) determined the charge on the electron in 1909.

17 Radioactivity Radioactivity is the spontaneous emission of radiation by an atom. It was first observed by Henri Becquerel. Marie and Pierre Curie also studied it. Three types of radiation were discovered by Rutherford:  particles  particles  rays

18 Discovery of the Nucleus
The Atom, circa 1900 The prevailing theory was that of the “plum pudding” model, put forward by Thomson. It featured a positive sphere of matter with negative electrons imbedded in it. Discovery of the Nucleus Ernest Rutherford shot  particles at a thin sheet of gold foil and observed the pattern of scatter of the particles.

19 The Nuclear Atom Since some particles were deflected at large angles, Thomson’s model could not be correct. Rutherford postulated a very small, dense nucleus with the electrons around the outside of the atom. Most of the volume of the atom is empty space.

20 Other Subatomic Particles
Protons were discovered by Rutherford in 1919. Neutrons were discovered by James Chadwick in 1932. Protons and electrons are the only particles that have a charge. Protons and neutrons have essentially the same mass. The mass of an electron is so small we ignore it.

21 Symbols of Elements Elements are symbolized by one or two letters. All atoms of the same element have the same number of protons, which is called the atomic number, Z. The mass of an atom in atomic mass units (amu) is the total number of protons and neutrons in the atom.

22 Isotopes Isotopes are atoms of the same element with different masses.
Isotopes have different numbers of neutrons.

23 © 2012 Pearson Education, Inc.
Atomic Mass Atomic and molecular masses can be measured with great accuracy using a mass spectrometer. Average Mass Because in the real world we use large amounts of atoms and molecules, we use average masses in calculations. Average mass is calculated from the isotopes of an element weighted by their relative abundances. © 2012 Pearson Education, Inc.

24 Periodic Table The periodic table is a systematic catalog of the elements. Elements are arranged in order of atomic number. The rows on the periodic chart are periods. Columns are groups. Elements in the same group have similar chemical properties.

25 Periodicity When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities. Metalloids border the stair-step line (with the exception of Al, Po, and At).

26 Groups These five groups are known by their names.
Nonmetals are on the right side of the periodic table (with the exception of H). Metals are on the left side of the chart.

27 Chemical Formulas The subscript to the right of the symbol of an element tells the number of atoms of that element in one molecule of the compound. Molecular compounds are composed of molecules and almost always contain only nonmetals.

28 Diatomic Molecules These seven elements occur naturally as molecules containing two atoms: Hydrogen Nitrogen Oxygen Fluorine Chlorine Bromine Iodine

29 Types of Formulas Empirical formulas give the lowest whole-number ratio of atoms of each element in a compound. Molecular formulas give the exact number of atoms of each element in a compound. Structural formulas show the order in which atoms are bonded. Perspective drawings show the three-dimensional array of atoms in a compound.

30 Ions When atoms lose or gain electrons, they become ions.
Cations are positive and are formed by elements on the left side of the periodic chart. Anions are negative and are formed by elements on the right side of the periodic chart.

31 Ionic Bonds Ionic compounds (such as NaCl) are generally formed between metals and nonmetals.

32 Writing Formulas Because compounds are electrically neutral, one can determine the formula of a compound this way: The charge on the cation becomes the subscript on the anion. The charge on the anion becomes the subscript on the cation. If these subscripts are not in the lowest whole-number ratio, divide them by the greatest common factor.

33 © 2012 Pearson Education, Inc.
Common Cations © 2012 Pearson Education, Inc.

34 © 2012 Pearson Education, Inc.
Common Anions © 2012 Pearson Education, Inc.

35 Inorganic Nomenclature
Write the name of the cation. If the anion is an element, change its ending to -ide; if the anion is a polyatomic ion, simply write the name of the polyatomic ion. If the cation can have more than one possible charge, write the charge as a Roman numeral in parentheses.

36 Patterns in Oxyanion Nomenclature
When there are two oxyanions involving the same element: The one with fewer oxygens ends in -ite. The one with more oxygens ends in -ate. NO2− : nitrite; SO32− : sulfite NO3− : nitrate; SO42− : sulfate Central atoms on the second row have bond to at most three oxygens; those on the third row take up to four. Charges increase as you go from right to left.

37 The one with the second fewest oxygens ends in -ite.
ClO2− : chlorite The one with the second most oxygens ends in -ate. ClO3− : chlorate The one with the fewest oxygens has the prefix hypo- and ends in -ite. ClO− : hypochlorite The one with the most oxygens has the prefix per- and ends in -ate. ClO4− : perchlorate

38 Acid Nomenclature If the anion in the acid ends in -ide, change the ending to -ic acid and add the prefix hydro- . HCl: hydrochloric acid HBr: hydrobromic acid HI: hydroiodic acid If the anion in the acid ends in -ite, change the ending to -ous acid. HClO: hypochlorous acid HClO2: chlorous acid If the anion in the acid ends in -ate, change the ending to -ic acid. HClO3: chloric acid HClO4: perchloric acid

39 Nomenclature of Binary Compounds
The less electronegative atom is usually listed first. A prefix is used to denote the number of atoms of each element in the compound (mono- is not used on the first element listed, however) .

40 © 2012 Pearson Education, Inc.
The ending on the more electronegative element is changed to -ide. CO2: carbon dioxide CCl4: carbon tetrachloride If the prefix ends with a or o and the name of the element begins with a vowel, the two successive vowels are often elided into one. N2O5: dinitrogen pentoxide © 2012 Pearson Education, Inc.

41 Nomenclature of Organic Compounds
Organic chemistry is the study of carbon. Organic chemistry has its own system of nomenclature. The simplest hydrocarbons (compounds containing only carbon and hydrogen) are alkanes. The first part of the names just listed correspond to the number of carbons (meth- = 1, eth- = 2, prop- = 3, etc.).

42 When a hydrogen in an alkane is replaced with something else (a functional group, like -OH in the compounds above), the name is derived from the name of the alkane. The ending denotes the type of compound. An alcohol ends in -ol.

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