Chapter 1 Learning Objectives - Atoms, Models, and Theories PDF

CLUE Chapter 1 Learning Objectives Chapter 1 Big Ideas Matter is made of atoms. Atoms are the smallest distinguishable part of an element. We use models (mental and physical) to represent many chemical entities. We use different models for different purposes. Theories (such as a...

CLUE Chapter 1 Learning Objectives Chapter 1 Big Ideas Matter is made of atoms. Atoms are the smallest distinguishable part of an element. We use models (mental and physical) to represent many chemical entities. We use different models for different purposes. Theories (such as atomic theory) change over time according to the evidence available. All atoms/molecules attract each other because of their attractive electrical forces. Attractions lower the potential energy of a system and repulsions tend to raise potential energy. Stable systems form where the attractive forces and the repulsive forces are equal. The temperature of a phase change depends on the strength of the intermolecular forces. Chapter 1 Learning Objectives 1.1 Make an argument (i.e. make a claim and support it with evidence and reasoning) for the existence of: a. Atoms - Claim: Atoms exist and make up all matter. Evidence: Dalton’s Atomic Theory (1803) proposed that matter comprises indivisible atoms. This idea was supported by the observation of chemical reactions, which show that elements combine in fixed ratios. Reasoning: The predictable behavior of substances in reactions and the consistency of mass conservation support the existence of atoms. b. Electrons - Claim: Electrons are negatively charged particles within atoms. - Evidence: J.J. Thomson’s cathode ray experiment (1897) discovered that cathode rays could be deflected by electric and magnetic fields, implying that they consisted of negatively charged particles (electrons). Reasoning: The behavior of cathode rays and precise measurements of electron charge strongly support the existence of electrons within atoms. c. The existence of a small, massive, and positively charged nucleus - Claim: Atoms have a small, dense, positively charged nucleus. Evidence: Rutherford’s gold foil experiment (1911) showed that most alpha particles passed through gold foil, but some were deflected. This suggested a tiny, dense, positively charged region (the nucleus) at the atom’s center. Reasoning: The deflection of particles in Rutherford’s experiment shows that most of the atom is empty space, with the dense, positively charged nucleus causing the observed deflections. 1.2 Draw diagrams/pictures of the various atom models as they changed over time. 1.3 Use the models to explain how and why the model of the atom changed over time as new evidence arose. - Dalton's Model (1803): John Dalton proposed that atoms were tiny, indivisible spheres. His model suggested that atoms of different elements had different weights, but he didn’t know about subatomic particles (like electrons or protons). This was the first step in atomic theory based on experimental evidence. - Thomson’s Model (1897): J.J. Thomson discovered the electron using experiments with cathode rays. This led him to propose the "plum pudding" model, where atoms  were made of a positively charged "pudding" with negatively charged electrons (like plums) scattered inside. This model suggested that atoms had internal structure, not just a solid sphere. - Rutherford’s Model (1911): Ernest Rutherford, through his famous gold foil experiment, showed that atoms have a tiny, dense nucleus at the center, which contains most of the atom's mass. The electrons were thought to orbit around this nucleus. This revealed that atoms were mostly empty space, with the nucleus being a small, dense core. - Bohr’s Model (1913): Niels Bohr built on Rutherford’s model by proposing that electrons travel in fixed orbits or energy levels around the nucleus. This explained why atoms didn’t collapse (as electrons wouldn’t spiral into the nucleus), and it also helped explain the specific energy levels of electrons observed in atomic spectra. 1.4 Explain how a scientific theory differs from everyday use of “theory”. - A Scientific theory is a comprehensive explanation based on extensive research and data, considered highly reliable within the scientific community while Everyday theory is a personal opinion or an untested idea, often used when someone is unsure or speculating about something. 1.5 Compare the parts of the various atomic theories that stayed the same over time and those that changed. - Parts that Stayed the Same: - Atoms as Basic Building Blocks: Since the time of Dalton (1803), all atomic theories agreed that atoms are the fundamental building blocks of matter. They are  indivisible (early theory) or made of smaller parts (later theories), but they are the foundation of all substances. - Atoms of Different Elements Are Different: Both Dalton and later scientists agreed that each element has unique atoms that differ from other elements in terms of mass and properties. - Parts that Changed: - Atom’s Structure: - Dalton's Model (1803): Atoms were solid, indivisible spheres. - Thomson’s Model (1897): Atoms were made of a positively charged "pudding" with electrons scattered like "plums." - Rutherford’s Model (1911): The atom has a tiny, dense, positively charged nucleus with electrons orbiting around it. - Bohr’s Model (1913): Electrons travel in fixed orbits around the nucleus, which explains atomic spectra. Subatomic Particles: Initially, Dalton’s theory didn’t recognize subatomic particles. Later, electrons, protons, and neutrons were discovered, adding complexity to our understanding of the atom’s structure. - Nucleus: Dalton’s theory had no concept of a nucleus. Rutherford proposed the nucleus, which was later refined by Bohr and others, who showed it contained protons and neutrons. 1.6 Develop a scientific question and a scientific explanation, and use evidence and data to make an argument. 1.7 Compare and contrast gravitational and electrostatic forces. - Electrostatic force: refers to the attractive or repulsive force that exists between two electrically charged objects. - Gravitational force: the attractive force that exists between any two objects with mass, essentially pulling them towards each other. - Similarities: -both meditated by fields -both proportional to 1/r² -both require 2 objects Differences: -gravitational is proportional to mass of objects -electrostatic is proportional to charge -gravitational is only attractive -electrostatic can be attractive or repulsive 1.8 Construct an atomic level explanation for why two isolated atoms would attract each other as they approach, and why they would repel if they get too close together. - When two isolated atoms approach each other, they initially experience an attractive force due to the electrostatic attraction between the positively charged nucleus of one atom and the negatively charged electron cloud of the other atom; however, if they get too close, the electron clouds start to overlap causing strong repulsion between the like- charged electrons, leading to a net repulsive force between the atoms.  1.9 Predict and explain the changes in the potential energy, the kinetic energy and the total energy as two isolated helium (or another noble gas) atoms approach each other. - As the He atoms approach the PE begins to decrease until the atoms reach PE well were they are a stable for a brief moment then the electron clouds overlap causing the PE to shoot up. Its the opposite for KE. total energy would stay the same because energy is neither created nor destroyed. - As the two atoms approach, the potential energy decreases (they get closer), the kinetic energy increases (they speed up), but the total energy stays the same (it’s conserved). - In other words, the potential energy goes up when the atoms are separated more, as they are less "held together" by the attractive forces between them. This is the opposite of when they approach each other and the potential energy decreases. 1.10 Draw an energy diagram showing potential energy as a function of internuclear distance. - 1.11 Construct a diagram/picture and use it to explain how energy is transferred at the atomic level (by collisions that can either add energy to the system or remove it). - When atoms or molecules collide, energy can either be transferred into the system (increasing the energy) or out of the system (decreasing the energy). 1. Adding Energy: If two atoms collide and one is moving faster, it can transfer energy to the other atom, making it move faster or vibrate more. This increases the kinetic energy of the atoms involved, which can cause them to heat up. 2. Removing Energy: On the other hand, if one atom is slower, it can lose energy in the collision, causing it to move slower or cool down. This transfer of energy can result in a decrease in the kinetic energy of the atom. 1.12 Predict/rank the relative London Dispersion Forces between atoms and molecules of different sizes. - Larger atoms or molecules with more electrons will have stronger London Dispersion Forces. So, if you rank them, larger molecules like iodine (I₂) or bigger hydrocarbons will have the strongest LDFs, while smaller atoms like helium or hydrogen will have the weakest. 1.13 Relate the strength of London Dispersion Forces to relative melting and boiling points for the noble gases or simple diatomic molecules (H2, N2, O2, F2, Cl2, Br2, and I2). - The stronger the London Dispersion Forces (due to increasing atomic size), the higher the melting and boiling points; larger atoms have more significant electron clouds, which are more easily polarized, leading to stronger temporary dipoles and stronger intermolecular attractions. - Polarized (mean?) - Large electron size = Large electron cloud = stronger dispersion forces = higher boiling points. - Helium has the lowest boiling point. Weak dispersion forces and smaller electron size. - The, smaller the electron cloud(the number of electrons= the number of protons(atomic number), the weaker the LDF the lower the boiling/melting point. Lowest to highest melting/boiling pt: H2,F2, N2, Cl2, Br2, I2, 02. - London Dispersion Forces are the only intermolecular force in nonpolar molecules like noble gases and most diatomic molecules. - The strength of London Dispersion Forces depends on the polarizability of the atoms, which increases with atomic size. - Higher polarizability leads to stronger temporary dipoles, resulting in stronger intermolecular attractions and higher boiling points. 1.14 Contrast the energy change when two helium or hydrogen atoms combine. Helium: - Helium is a noble gas. And all noble gases have their electron shells filled. If their shells are already filled, they do not need another bond to fill the electron shells. So, they don’t readily bond. (Remember that bonds happen because there is low or high electrons in the electron shells). Energy changes only when there is an electron transfer in the electron shells. - Helium atoms have a full valence shell, so when they come close, there is no opportunity for significant electron sharing, leading to minimal interaction and essentially no energy change upon "combination". - When two helium atoms combine, there is essentially no energy change as they do not readily form a bond due to their already filled electron shells, meaning the  interaction between them is very weak and only involves weak van der Waals forces; essentially, no significant energy is released or absorbed. Hydrogen: - When two hydrogen atoms approach each other, their electron clouds overlap, allowing them to share electrons and form a covalent bond, resulting in a decrease in potential energy and energy release. 1.15 Differentiate between London Dispersion Forces and covalent bonding without providing an explanation. - London Dispersion Forces are significantly weaker than covalent bonds. - London Dispersion Forces occur between molecules (intermolecular), while covalent bonds occur within a molecule (intramolecular). - London Dispersion Forces arise from temporary fluctuations in electron distribution, while covalent bonds result from sharing of electrons between atoms. - London Dispersion Forces act between molecules, while covalent bonds hold atoms together within a molecule.