A molecular approach 1-4 Chapter summary PDF

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Summary

This is a summary of chapters 1-4, which summarizes:
Chapter 1. Matter, Measurement, and Problem Solving
Chapter 2. Atoms and Elements
Chapter 3. Molecules, Compounds, and Chemical Equations
Chapter 4. Chemical Quantities and Aqueous Reactions...

Description

A molecular approach: Chapter 1-4 Note/Summary Chapter 1: Matter, Measurement, and Problem Solving  Carbon monoxide gas is collected of carbon monoxide molecules.  Each molecule covers a carbon atom and an oxygen atom held together by a chemical bond.  Atoms are the submicroscopic atoms that constitute the fundamental building blocks of ordinary matter.  Free atoms are infrequent in nature; in its place they bind together in specific geometrical

arrangements to form molecules.  Chemistry is the science that seeks to comprehend the behavior of matter by studying the behavior of atoms and molecules.

 Chemistry is the science that seeks to understand the behavior of matter by studying the behavior of atoms and molecules. Section 1.1 the Scientific Approach to Knowledge  The scientific method is a process for understanding nature by observing nature and its behavior, and by conducting experiments to test our ideas.  Scientific method include observation, formulation of hypotheses, experimentation, and formulation of laws and theories.  Observations often lead scientists to formulate a hypothesis. The Scientific Method OBSERVATION

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EXPLANATION

The properties of matter are determined by the atoms and molecules that compose it Carbon Monoxide  composed of one carbon atom and one oxygen atom  colorless, odorless gas  burns with a blue flame  binds to hemoglobin

Carbon Dioxide  composed of one carbon atom and two oxygen atoms  colorless, odorless gas  incombustible  does not bind to hemoglobin

Section 1.2 the Scientific Approach to Knowledge  philosophers attempt to comprehend the universe by reasoning and thinking about “ideal” behavior  Scientists willing to understand the universe through  Empirical information gained finished observation and experiment Hypothesis Falsifiable – confirmed or refuted by other observations Tested by experiments – validated or invalidated When similar observations are consistently made, it can lead to a Scientific Law  a statement of a behavior that is always observed  summarizes past observations and predicts future ones

 Law of Conservation of Mass A theory is a general explanation for the manifestation and behavior of all nature  Models  Pinnacles Of Scientific knowledge  validated or invalidated by experiment and observation Figure of Scientific

Section 1.3 Classification of Matter  Matter is anything that has mass and occupies space  We can classify matter based on whether it’s solid, liquid, or gas

 matter can be classified as solid, liquid, or gas based on the characteristics Solid Shape and volume is fixed

There are three type of state solid, Liquid and Gas whereas shape fixed, Indefinite and Indefinite and volume is Fixed, fixed and Indianite.  Fixed = keeps shape when placed in a container  Indefinite = takes the shape of the container Solids  the particles in a solid are packed close together and are fixed in position  the close packing of the particles results in solids being incompressible  The inability of the particles to move around results in solids retaining their shape and volume when placed in a new container, and prevents the particles from flowing

Crystalline Solids  some solids have their particles arranged in an orderly geometric pattern – we call these crystalline solids E.g Salt and diamonds

Liquid

 The particles in a liquid are closely packed, but they have some ability to move around  The close packing results in liquids being incompressible  But the ability of the particles to move allows liquids to take the shape of their container and to flow. However, they don’t have enough freedom to escape and expand to fill the container

Gases  In the gas state, the particles have complete freedom from each other  The particles are constantly flying around, bumping into each other and the container  In the gas state, there is a lot of empty space between the particles

Gases  because there is a lot of empty space, the particles can be squeezed closer together – therefore gases are compressible  Because the particles are not held in close contact and are moving freely, gases expand to fill and take the shape of their container, and will flow

Section 1.4 Classification of Matter by Composition  Matter whose composition does not change from one sample to another is called a pure substance  Because composition is always the same, all samples have the same characteristics  Matter whose composition may vary from one sample to another is called a mixture

Classification of Matter by Composition

Pure Substance  made of one type of particle  all samples show the same intensive properties Mixture  made of multiple types of particles  samples may show different intensive properties Classification of Pure Substances

 Substances that cannot be broken down into simpler substances by chemical reactions are called elements  Substances that can be decomposed are called compounds  Most natural pure substances are compounds Classification of Pure Substances  Made of one type of atom (some elements found as multi-atom molecules in nature)  Combine together to make compounds

 Made of one type of molecule, or array of ions  Molecules contain 2 or more different kinds of atoms Classification of Mixtures  Made of multiple substances, whose presence can be seen  Portions of a sample have different composition and properties

 Made of multiple substances, but appears to be one substance  All portions of a sample have the same composition and properties Section 1.5 Changes in Matter  Changes that alter the state or appearance of the matter without altering the composition are called physical changes  Changes that alter the composition of the matter are called chemical changes Physical Changes in Matter

The boiling of water is a physical change. The water molecules are separated from each other, but their structure and composition do not change. Chemical Changes in Matter

The rusting of iron is a chemical change. The iron atoms in the nail combine with oxygen atoms from O2 in the air to make a new substance, rust, with a different composition. Properties of Matter  Physical properties are the characteristics of matter that can be changed without changing its composition  Chemical properties are the characteristics that determine how the composition of matter changes as a result of contact with other matter or the influence of energy Section 1.6 Solving Chemical Problems Units  Always write every number with its associated unit  Always include units in your calculations You can do the same kind of operations on units as you can with numbers cm × cm = cm2 cm + cm = cm cm ÷ cm = 1 Using units as a guide to problem solving is called dimensional analysis

 Many problems in chemistry involve using relationships to convert one unit of measurement to another  Conversion factors are relationships between two units  Conversion factors generated from equivalence statements  Arrange conversion factors so given unit cancels  May string conversion factors Conceptual Plan  a conceptual plan is a visual outline that shows the strategic route required to solve a problem  for unit conversion, the conceptual plan focuses on units and how to convert one to another  for problems that require equations, the conceptual plan focuses on solving the equation to find an unknown value How many cubic centimeters are there in 2.11 yd3? Can use density as a conversion factor between mass and volume!! Density of H2O = 1.0 g/mL \ 1.0 g H2O = 1 mL H2O Density of Pb = 11.3 g/cm3 \ 11.3 g Pb = 1 cm3 Pb Problem Solving  When solving a problem involves using an equation, the concept plan involves being given all the variables except the one you want to find  Solve the equation for the variable you wish to find, then substitute and compute

Chapter 2 Atoms and Elements  Each element has a unique number of protons in its nucleus  The number of protons in the nucleus of an atom is called the atomic number Electrons The number of electrons (-) must equal the number of protons (+) in a neutral atom. Atomic Mass  The atomic mass is the number of protons and neutrons.  In order to determine the number of neutrons, subtract the atomic number from the atomic mass.

Section 2.1 Structure of the Nucleus  Soddy discovered that the same element could have atoms with different masses, which he called isotopes 

there are two isotopes of chlorine found in nature, one that has a mass of about 35 amu and another that weighs about 37 amu

 The observed mass is a weighted average of the weights of all the naturally occurring atoms



the percentage of an element that is one isotope is called the isotope’s percent abundance



If the atomic mass of chlorine is 35.45 amu, which isotope is more abundant, Cl-35 or Cl-37?

Isotopes  All isotopes of an element are chemically identical  All isotopes of an element have the same number of protons  Isotopes of an element have different masses  Isotopes of an element have different numbers of neutrons  Isotopes are identified by their mass numbers, which is the sum of all the protons and neutrons in the nucleus

Neon

Number of Portion

Number of Neutrons

Percent Natural Abundance Symbol Ne-20

10

10

10

11

90.48% Ne-21 0.27%

Ne-22

10

12

9.25%

Section 2.2 Charged Atoms  When atoms gain or lose electrons, they acquire a charge  Charged atoms or groups of atoms are called ions  When atoms gain electrons, they become negatively charged ions, called anions  When atoms lose electrons, they become positively charged ions, called captions Ions and Compounds  Ions behave much differently than the neutral atoms 

the metal sodium, made of neutral Na atoms, is highly reactive and quite unstable; however, the sodium cations, Na+, found in table salt are very nonreactive and stable

 Because materials such as table salt are neutral, there must be equal amounts of charge from cations and anions in them Metals  Solids at room temperature, except Hg  Reflective surface



Shiny

 Conduct heat  Conduct electricity  Malleable  Ductile  Lose electrons and form cations in reactions  About 75% of the elements are metals  Left side of table (except H)

Nonmetals  Found in all three states  Poor conductors of heat  Poor conductors of electricity  Solids are brittle  Gain electrons in reactions to become anions  Right on the table

Metalloids  Show some properties of metals and some of nonmetals  Also known as semiconductors

Section 2.3 the Modern Periodic Table  Elements with similar chemical and physical properties are in the same column  Columns are called Groups or Families  Rows are called Periods  Each period shows the pattern of properties repeated in the next period Important Groups – Alkali Metals

 Group I = Alkali Metals  Hydrogen usually placed here, though it doesn’t really belong  Soft, low melting points, low density

 Flame tests ® Li = red, Na = yellow, K = violet  Very reactive, never find uncombined in nature  Group II = Alkali earth metals  Harder, higher melting, and denser than alkali metals 

Mg alloys used as structural materials

 Flame tests ® Ca = red, Sr = red, Ba = green  Reactive, but less than corresponding alkali metal Important Groups – Halogens  Group 7 = halogens  Nonmetals  F2 and Cl2 gases; Br2 liquid; I2 solid  All diatomic  Very reactive

Important Groups – Noble Gases  Group 8= Noble Gases  All gases at room temperature  Very unreactive, practically inert  Very hard to remove electron from or give electron to

Section 2.4 Ion Charge and the Periodic Table  The charge on an ion can often be determined from an element’s position on the Periodic Table  Metals always form positively charged cations  For many main group metals, the charge = the group number  Nonmetals form negatively charged anions  For nonmetals, the charge = the group number – 8 What is the charge on each of the following ions?  potassium cation

K+

 sulfide anion

S2−

 calcium cation

Ca2+

 bromide anion

Br−

 aluminum cation

Al3+

Atoms by Moles  If we can find the mass of a particular number of atoms, we can use this information to convert the mass of an element sample into the number of atoms in the sample  The number of atoms we will use is 6.022 x 1023, and we call this a mole 1 mole = 6.022 x 1023 things like 1 dozen = 12 things

Chemical Packages – Moles  The number of particles in 1 mole is called Avogadro’s Number = 6.0221421 x 1023 

A particle can be anything: atom, molecule, ion, electron, proton, etc.

Section 2.5 Relationship between Moles and Mass  The mass of one mole of atoms is called the molar mass  The molar mass of an element, in grams, is numerically equal to the element’s atomic mass, in amu

Section 2.6 Problem & Solution Which of the following represents the least number of molecules? A. 20.0 g of H2O (18.02 g/mol) B. 77.0 g of CH4 (16.06 g/mol) C. 68.0 g of CaH2 (42.09 g/mol) D. 100.0 g of N2O (44.02 g/mol) E. 84.0 g of HF (20.01 g/mol)

Solution 20.0 g of H2O represents the smallest number of moles, meaning the least number of molecules present. Since 1 mole = 6.022 × 1023 molecules (or atoms) regardless of identity, the least number of moles will equal the least number of molecules. Problem 2: Compare 1 mole of H2, 1 mole of O2, and 1 mole of F (a) Which has the largest number of molecules? Explain why. (b) Which has the greatest mass? Explain why. Answer a 1 mole is always 6.022 x 1023 molecules. They have the same number of molecules. Answer b F2; it has the highest molar mass.

Chapter 3: Compounds & Molecules  Compound is a combination of 2 or more elements in definite ratios by mass.  The character of each element is lost when forming a compound (e.g., think of NaCl).  Molecules are the smallest units of a compound that retains the characteristics of the compound. Section 3.1 Formulas

 Formula for glycine is C2H5NO2  In one molecule there are 

2 C atoms



5 H atoms



1 N atom



2 O atoms

Condensed Formula  Formula for glycine is NH2CH2CO2H  In one molecule there are 

1 NH2 (amine group)



1 CH2 group



1 CO2H group

Structural Formula  Show how the atoms are attached within a molecule  The lines between atoms represent chemical bonds that hold the atoms together.

WRITING FORMULAS  Can also write glycine as the condensed formula H2NCH2COOH  Or in the form of a structural formula

Section 3.2 Resources for Molecular Modeling  Modeling software 

CAChe (General Chemistry Interactive CD-ROM)



Rasmol



Molden



Gaussview



Maestro

Molecular weight = sum of the atomic weights of all atoms in the molecule. Molar mass = molecular weight in grams per mol. What is the molar mass of ethanol, C2H6O? 1 mol contains 2 moles of C (12.01 g C/1 mol) = 24.02 g C 6 moles of H (1.01 g H/1 mol) = 6.06 g H 1 mol of O (16.00 g O/1 mol) = 16.00 g O TOTAL = molar mass = 46.08 g/mol How many moles of alcohol (C2H6O) are there in a “standard” can of beer if there are 21.3 g of C2H6O?  Molar mass of C2H6O = 46.08 g/mol

 Calc. moles of alcohol Molecular & Ionic Compounds

Molecular compounds consist of discrete molecules Ionic compounds consist of discrete ions Ions & Ions Compounds  IONS are atoms or groups of atoms with a positive or negative charge.  Taking away an electron from an atom gives a CATION with a positive charge  Adding an electron to an atom gives an ANION with a negative charge. Section 3.3 Predicting Ion Charges  metals (Mg) lose electrons ---> cations  nonmetals (F) gain electrons ---> anions Metals M ---> n e- + Mn+ Where n = periodic group Na+

sodium ion

Mg2+

magnesium ion

Al3+

aluminum ion

Transition metals --> M2+ or M3+ Are common

Fe2+

iron(II) ion

Fe3+

iron(III) ion

Ion Formation

Reaction of aluminum and bromine

Polyatomic Ions  Groups of atoms with a charge.  MEMORIZE the names and formulas of common  NH4+ ammonium ion  One of the few common polyatomic cations Polyatomic Ions (oxyanions)

 Prefix per- and suffix –ate: largest #

 Suffix -ate : greater # of oxygen atoms  Suffix -ite : smaller # of oxygen atoms  Prefix hypo- and suffix –ite: smallest #

 SO42- sulfate ion  SO32- sulfite ion

 NO3- nitrate ion  NO2- nitrite ion Section 3.4 Compounds and formed from Ions  Cation + Anion --->  COMPOUND  Na+ + Cl- --> NaCI  A neutral compd. requires equal number of + and - charges.

Some Ionic Compounds  Ca2+ + 2 F- ---> CaF2  Mg2+ + NO3- ----> Mg(NO3)2  Fe2+ + PO43- ----> Fe3(PO4)2

Properties of Ionic Compounds Forming Naci from Na and CI2  A metal atom can transfer an electron to a nonmetal.  The resulting cation and anion are attracted to each other by electrostatic forces.

Electrostatic Forces

The oppositely charged ions in ionic compounds are attracted to one another by ELECTROSTATIC FORCES. These forces are governed by COULOMB’S LAW. Importance of Coulomb’s Law

NaCl, Na+ and Cl-, m.p. 804 oC

Molecular Compounds

CO2 Carbon dioxide

BCI3 boron trichloride

MgO, Mg2+ and O2- m.p. 2800 oC

Naming Molecular Compounds  All are formed from two or more nonmetals.  Ionic compounds generally involve a metal and nonmetal (NaCl)

Empirical & Molecular Formulas  A pure compound always consists of the same elements combined in the same proportions by weight.  Therefore, we can express molecular composition as percent by weight 

Ethanol, C2H6O



52.13% C



13.15% H



34.72% O

Percent Composition  Consider some of the family of nitrogen-oxygen compounds: 

NO2, nitrogen dioxide and closely related, NO, nitrogen monoxide (or nitric oxide)

Section 3.5 Determining Formulas

 In chemical analysis we determine the % by weight of each element in a given amount of pure compound and derive the EMPIRICAL or SIMPLEST formula. A compound of B and H is 81.10% B. What is its empirical formula?

 Because it contains only B and H, it must contain 18.90% H.  In 100.0 g of the compound there are 81.10 g of B and 18.90 g of H.  Calculate the number of moles of each constituent. A compound of B and H is 81.10% B. Its empirical formula is B2H5. What is its molecular formula? Is the molecular formula B2H5, B4H10, B6H15, B8H20, etc.?

Section 3.6 Problem & Solution

For each of the following compounds, state whether it is ionic or covalent. If it is ionic, write the symbols for the ions involved: (a) NF3 (b) BaO (c) (NH4)2CO3 (d) Sr(H2PO4)2 (e) IBr (f) Na2O Solution A. Covalent B. ionic, Ba2+, O2− C. ionic, NH+4, CO2− D. ionic, Sr2+, H2PO− E. covalent F. ionic, Na+, O2−

Chapter 4 Chemical Quantities and...