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Download Integrated Science Chapter 3 Notes Section 1: Atomic Structure 1
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Integrated Science Chapter 3 Notes Section 1: Atomic Structure 1. What are Atoms • Around 400 B.C. the Greek philosopher Democritius suggested the universe was composed of invisible units called atoms The word atom is derived from the Greek word meaning “unable to be divided” Democritius said there were only four types of atoms: air, water, earth, and fire The problem with Democritius’ idea was that it could not be proven • In 1808, an English schoolteacher named John Dalton proposed an atomic theory Dalton’s theory had three main ideas: ♦ Every element is made of tin, unique particles called atoms that cannot be divided ♦ Atoms of the same element are exactly alike ♦ Atoms of different elements can join to form molecules Dalton’s theory, like Democritius’, could not be proven • An atom is the smallest part of an element that still has the element’s properties 2. What’s in an Atom • Less than 100 years after Dalton published his atomic theory, scientists determined that atoms could be split. • There are three subatomic particles involved in the everyday chemistry of most substances Proton – a positively charged subatomic particle in the nucleus of an atom Neutron – a neutral subatomic particle in the nucleus of an atom ♦ Nucleus – the center of an atom; made up of protons and neutrons Electron – a tiny negatively charged subatomic particle moving around outside the nucleus of an atom • Although the atom is composed of both charged particles and neutral particles, the atom has no overall charge The number of protons (+ charge) equals the number of electrons (- charge) Particle Proton Charge +1 Mass (kg) 1.67 x 10-27 Location in atom In the nucleus Neutron 0 1.67 x 10-27 In the nucleus -1 -31 Electron 9.11 x 10 Outside the nucleus 3. Models of the Atom • In 1913, the Danish scientist Niels Bohr suggested that electrons in an atom move in set paths around the nucleus Each electron has a certain energy that is determined by its path. The path is called an energy level Energy level – any of the possible energies an electron may have in an atom • By 1925, Bohr’s model of the atom no longer explained all observations A new atomic model proposed that electrons behave more like waves on a vibrating string than like particles In the modern model of the atom it is impossible to determine the exact location of an electron that is moving around the nucleus of an atom, scientists could only calculate the chance of finding an electron in a certain place with in the atom Orbital – a region in an atom where there is a high probability of finding electrons ♦ The orbitals have several different orientations in space ♦ There are four different types of orbitals: s, p, d, and f. The s orbital is the lowest energy orbital, and the f orbital is the highest energy orbital. ♦ Electrons fill the orbitals from lowest energy to highest energy, and each orbital can only contain two electrons Valence electron – an electron in the outermost energy level of an atom Section 2: The Periodic Table 1. The periodic table groups similar elements together • The elements are also arranged in a certain order based on the number of protons an atom of that element has in its nucleus • Periodic law – properties of elements tend to change in a regular pattern when elements are arranged in order of increasing atomic number or the number of protons in their atoms periods - a horizontal row of elements in the periodic table ♦ valence electrons determine the chemical properties of atoms group – a vertical column of elements in the periodic table ♦ atoms of electrons in the same group have the same number of valence electrons, so these elements have similar properties 2. Some Atoms Form Ions • ionization – the process of addling electrons to or removing electrons from at atom or groups of atoms atoms will gain or lose electrons I order to have a full outermost energy level ♦ by gaining or losing electrons the atom no longer has a neutral charge, it becomes a charged particle ion – an atom or group of atoms that has lost or gained one or more electrons and therefore has a net electric charge ♦ cation – an ion with a positive charge ♦ anion – an atom with a negative charge 3. The Structure of Atoms • Because atoms have different structures, they have different properties • Atomic number – the number of protons in the nucleus of an atom • Mass number – the total number of protons and neutrons in the nucleus of an atom • Isotopes – any atoms having the same number of protons but different numbers of neutrons • The mass of a single atom is very small Example: the mass of a single fluorine atom is less than one trillionth of a billionth of a gram (1/1,000,000,000,000,000,000,000 g) The mass of on an atom is expressed in atomic mass units Atomic mass unit (amu) – a quantity equal to one-twelfth of the mass of a carbon-12 atom (the mass of one hydrogen atom is approximately 1 amu) When compared to the proton and neutron, the electron has very little mass (it would take about 2,000 electrons to have the same mass as one proton) Average atomic mass – the weighted average of the masses of all naturally occurring isotopes of an element Section 3: Families of Elements 1. groups in the periodic table are sometimes called families • elements in the same family have similar properties 2. elements can be classified based on shared properties • metals – the elements that are good conductors of heat and electricity • nonmetals – the elements that are usually poor conductors of heat and electricity • semiconductors – the elements that are intermediate conductors of heat and electricity semiconductors are also called metalloids 3. most of the elements are classified as metals • there are four families of metals alkali metals – the highly reactive metallic elements located in Group 1 of the periodic table ♦ the alkali metals are very reactive because it has one valence electron that can be easily removed to forma a positive ion. Alkaline earth metals – the reactive metallic elements located in Group 2 of the periodic table ♦ The alkaline earth metals have two valence electrons, so are less reactive alkali metals Transition metals – the metallic elements located in Groups 3 – 12 of the periodic table ♦ Transition metals are less reactive than the alkaline earth metals ♦ Like the alkali and alkaline earth metals, they lose electrons to form positive ions, but the transition metals can lose different numbers of electrons 4. Nonmetals and Semiconductors • Nonmetals and their compound are plentiful on Earth Halogens – the highly reactive elements located in Group 17 of the periodic table ♦ The halogen gases need to acquire only one electron to have a full outermost energy level, that is why they are so reactive Noble gases – the unreactive gaseous elements located in Group 18 of the periodic table ♦ The noble gases exist as a single atoms instead of as molecules • The semiconductors will conduct heat and electricity only under certain conditions This property makes them useful in computer applications Section 4: Using Moles to Count Atoms 1. When people count out large numbers of small things, they often simply the job by using counting units • Counting units can be based on the size of a container or an exact number example: the counting unit for buying popcorn at the theater is the size of the container, small, medium, or large example: the counting unit for eggs is the dozen, or exactly 12 eggs • Because chemists often deal with large numbers of small particle they use a counting unit called the mole mole – the SI base unit that describes the amount of a substance One mole equals 602,213,670,000,000,000,000,000 particles, or 6.022 x 1023 particles This number is known as Avogadro’s constant and is the number of particles in 1 mole Molar mass – the mass in grams of one mole of a substance 2. Calculating moles • Because the amount of a substance and its mass are related, it is often useful to convert moles to grams, and grams to moles. An element’s molar mass acts as a conversion factor between moles and mass example: what is the mass in grams of 2 mols of hydrogen? 2.0 mol H x 1.008 gH = 2.016 g H mol H example: how many moles of oxygen in 5 g of oxygen? 5.0 g O x 1 mol O = 0.31 mol O 16.0 g O
